JOHN  ALEXANDER  JAMESON,  JR. 
1903-1934 


ENGINEERING  LIBRARY 


THIS  BOOK  belonged  to  John  Alexander  Jameson,  Jr.,  A.B.,  Wil- 
liams, 1925;  B.S.,  Massachusetts  Institute  of  Technology,  1928; 
M.S.,  California,  1933.  He  was  a  member  of  Phi  Beta  Kappa,  Tau 
Beta  Pi,  the  American  Society  of  Civil  Engineers,  and  the  Sigma 
Phi  Fraternity.  His  untimely  death  cut  short  a  promising  career. 
He  was  engaged,  as  Research  Assistant  in  Mechanical  Engineering, 
upon  the  design  and  construction  of  the  U.  S.  Tidal  Model  Labora- 
tory of  the  University  of  California. 

His  genial  nature  and  unostentatious  effectiveness  were  founded 
on  integrity,  loyalty,  and  devotion.  These  qualities,  recognized  by 
everyone,  make  his  life  a  continuing  beneficence.  Memory  of  him 
will  not  fail  among  those  who  knew  him. 


' 


INTERNATIONAL   CHEMICAL  SERIES 
H.  P.  TALBOT,  PH.D.,  Sc.D.,  CONSULTING  EDITOR 


A  TEXTBOOK 

OF 

INORGANIC  CHEMISTRY 
FOR  COLLEGES 


INTERNATIONAL  CHEMICAL  SERIES 
(H.  P.  TALBOT,  PH.D.,  Sc.D.,  CONSULTING  EDITOR) 


Bancroft — 

APPLIED   COLLOID   CHEM- 
ISTRY 
Bingha  m— 

FLUIDITY  AND  PLASTICITY 
Cady— 

INORGANIC  CHEMISTRY 
Cady — 

GENERAL  CHEMISTRY 
Griffin— 

TECHNICAL   METHODS   OF 
ANALYSIS 

As  Employed  in  the  Labora- 
tories of  Arthur  D.  Little,  Inc. 
Hall  and   W  illiams— 

CHEMICAL  AND  METALLO- 
GRAPHIC    EXAMINATION 
OF     IRON,     STEEL     AND 
BRASS 
Hamilton  and  Simpson — 

CALCULATIONS    OF    QUAN- 
TITATIVE 
ANALYSIS 
Loeb— 

PROTEINS         AND         THE 
THEORY   OF   COLLOIDAL 
BEHAVIOR 
Lord  and  Demorest — 

METALLURGICAL      ANALY- 
SIS 

Fifth  Edition 
Mahin— 

QUANTITATIVE  ANALYSIS 
Third  Edition 
Mahin  and  Carr — 

QUANTITATIVE     AGRICUL- 
TURAL ANALYSIS 
Millard— 

PHYSICAL  CHEMISTRY  FOR 
COLLEGES 


CHEMICAL 


Moore — 

HISTORY    OF    CHEMISTRY 
Norris — 

TEXTBOOK  OF  INORGANIC 
CHEMISTRY       FOR    COL- 
LEGES 
Norris  and  Mark — 

LABORATORY      EXERCISES 
IN   INORGANIC   CHEMIS- 
TRY 
Norris — 

ORGANIC  CHEMISTRY 
Second  Edition 
Norris — 

EXPERIMENTAL     ORGANIC 

CHEMISTRY 
Second  Edition 
Parr— 

ANALYSIS    OF    FUEL,    GAS. 
WATER  AND  LUBRICANTS 
Third  Edition 
Robinson — 

THE  ELEMENTS   OF  FRAC- 
TIONAL DISTILLATION 
White— 
TECHNICAL  GAS  AND  FUEL 

ANALYSIS 
Second  Edition 
Williams — 

PRINCIPLES  OF  METALLO- 
GRAPHY 
Woodman — 

FOOD  ANALYSIS 
Second  Edition 
Long  and  Anderson — 

CHEMICAL  CALCULATIONS 
BoQue — 

THE    THEORY  AND  APPLI- 
CATION    OF     COLLOIDAL 
BEHAVIOR 
Two  Volumes 


A  TEXTBOOK 

OF 

INORGANIC  CHEMISTRY 
FOR  COLLEGES 


BY 

JAMES  F.  NORRIS 

PROFESSOR  OF  ORGANIC  CHEMISTRY 

MASSACHUSETTS  INSTITUTE  OF  TECHNOLOGY 

AUTHOR  OF  "THE  PRINCIPLES  OF  ORGANIC   CHEMISTRY," 

"EXPERIMENTAL  ORGANIC  CHEMISTRY,"  AND   (WITH  K.  L.  MARK) 

"LABORATORY  EXERCISES  IN  INORGANIC  CHEMISTRY" 


FIRST  EDITION 
FIFTH  IMPRESSION 


McGRAW-HILL  BOOK  COMPANY,  INC. 

NEW  YORK:  370  SEVENTH  AVENUE 

LONDON:  6  &  8  BOUVERIE  ST.,  E.  C.  4 

1921 


Xt, 


COPYRIGHT,  1921,  BY  THE 
MCGRAW-HILL  BOOK  COMPANY,  INC. 


PRINTKQ.IN^THE    UNITED,  STATES    OF    AMERICA 


PREFACE 


FOR  a  number  of  years  the  author  of  this  textbook  had  the 
opportunity  to  teach  students  who  were  beginning  the  study  of 
chemistry.  His  experience,  acquired  in  the  recitation  room  and 
in  the  laboratory,  led  him  to  the  view  that  the  average  student 
finds  it  difficult  to  understand  many  of  the  apparently  simple 
concepts  of  the  science.  It  appeared,  therefore,  to  be  an  interest- 
ing task  to  attempt  to  present  the  material  commonly  treated  in 
elementary  books  on  chemistry  in  a  form  which  could  be  reason- 
ably well  followed  by  the  student  through  private  study  and  with 
the  smallest  amount  of  explanation  on  the  part  of  the  teacher. 
Since  this  book  has  been  written  from  this  point  of  view,  the  sub- 
ject has  been  developed  slowly,  and  the  consideration  of  the  more 
abstruse  material  has  been  deferred  until  the  student  has  gained 
some  familiarity  with  chemical  phenomena  and  with  the  language 
of  the  science.  No  attempt  has  been  made  at  conciseness  in  the 
discussion  of  important  principles.  Analogies  have  been  repeat- 
edly pointed  out  in  an  endeavor  to  indicate  to  the  student  the  way 
in  which  he  should  classify  the  facts  brought  to  his  attention. 

The  aim  of  the  author  has  been  to  present  the  general  princi- 
ples underlying  the  science;  as  a  consequence,  chemical  phe- 
nomena have  been  discussed  from  the  standpoints  of  both  matter 
and  energy.  The  law  of  mobile  equilibrium  in  its  broadest  sense 
has  been  used  repeatedly  in  interpreting  many  important  facts. 
The  more  elementary  parts  of  thermochemistry  and  electrochemis- 
try have  also  been  emphasized.  In  fact,  physical  chemistry  has 
been  drawn  on  frequently,  but  an  endeavor  has  been  made  to 
limit  its  use  to  the  elucidation  of  the  more  important  facts  of 
inorganic  chemistry. 

Several  chapters  of  the  book  are  devoted  to  the  consideration, 
in  a  general  way,  of  the  physical  and  chemical  properties  of  metals, 
non-metals,  acids,  bases,  and  salts.  It  is  possible  by  this  pro- 


vi  PREFACE 

cedure  to  bring  out  generalizations  of  value  and  to  emphasize  the 
relations  that  exist  between  the  uses  of  compounds  and  their 
properties.  Important  descriptive  matter  has  not  been  omitted, 
and  the  more  recent  advances  in  chemistry  in  both  its  technical 
and  theoretical  aspects  have  been  included.  The  modern  con- 
ception of  the  atom  has  been  presented  in  an  elementary  way. 
Throughout  the  book  an  attempt  has  been  made  to  impress  upon 
the  student  the  fact  that  chemistry  is  a  growing  science. 

A  large  number  of  exercises  are  placed  at  the  ends  of  most  of 
the  chapters.  In  the  main,  these  are  not  direct  questions  on  the 
text,  but  are  designed  to  furnish  the  student  with  an  opportunity 
to  use  his  knowledge.  Some  of  the  questions  may  tax  the  abil- 
ity of  even  the  best  students;  these  may  serve  as  the  basis  for 
discussions  in  the  classroom. 

The  author  is  greatly  indebted  to  his  wife  for  the  efficient 
help  which  he  has  received  throughout  all  the  work  required  in 
the  preparation  of  the  book.  This  assistance  included  the  copy- 
ing of  the  manuscript  and  help  in  proofreading  and  in  the  prepara- 
tion of  the  index.  He  is  also  indebted  to  Professor  K.  L.  Mark, 
of  Simmons  College,  for  helpful  suggestions  and  criticisms  based 
on  a  careful  reading  of  the  manuscript. 

Four  of  the  cuts  have  been  borrowed,  with  the  permission  of 
the  publishers,  from  Cady's  Inorganic  Chemistry.  These  include 
the  excellent  drawings  to  represent  the  chamber  process  for  sul- 
phuric acid  and  the  changes  which  take  place  in  a  blast  furnace. 
Other  cuts  have  been  adapted  from  Leighou's  Chemistry  of 
Materials  and  Black  and  Conant's  Practical  Chemistry. 

JAMES  F.  NORMS. 

CAMBRIDGE,  MASS. 
May  19,  1921. 


CONTENTS 


CHAPTER  PAGE 

I.  INTRODUCTION 1 

II.  PHYSICAL  AND  CHEMICAL  CHANGES 6 

III.  ELEMENTS  AND  COMPOUNDS 15 

IV.  OXYGEN 21 

V.  HYDROGEN 38 

VI.  THE  ATOMIC  THEORY.     CHEMICAL  EQUATIONS 52 

VII.  CHEMICAL  CALCULATIONS 68 

VIII.  MEASUREMENT  OF  GASES 77 

IX.  WATER 88 

X.  CHLORINE.     VALENCE 99 

XI.  HYDROCHLORIC  ACID.     DOUBLE  DECOMPOSITION 123 

XII.  THE  ENERGY  FACTOR  IN  CHEMICAL  CHANGE 137 

XIII.  OZONE  AND  HYDROGEN  PEROXIDE 144 

XIV.  PROPERTIES  OF  GASES,  LIQUIDS,  AND  SOLIDS 155 

XV.  CARBON  AND  ITS  OXIDES 171 

XVI.  COAL,  COKE,  ILLUMINATING  GAS,  FLAMES 194 

XVII.  ACIDS,  BASES,  SALTS.     SOLUTIONS 209 

XVIII.  CHEMICAL  EQUILIBRIUM 233 

XIX.  SULPHUR  AND  HYDROGEN  SULPHIDE 245 

XX.  THE  OXIDES  AND  ACIDS  OF  SULPHUR 256 

XXI.  NITROGEN  AND  THE  ATMOSPHERE 284 

XXII.  AMMONIA  AND  ITS  DERIVATIVES 302 

XXIII.  NITRIC  ACID,  NITROUS  ACID,  AND  THE  OXIDES  OF  NITROGEN.  320 

XXIV.  THE  DETERMINATION  OF  ATOMIC  AND  MOLECULAR  WEIGHTS.  345 
XXV.  THE  PERIODIC  LAW 358 

XXVI.  THE  HALOGEN  FAMILY 364 

XXVII.  SELENIUM  AND  TELLURIUM 392 

XXVIII.  PHOSPHORUS,   ARSENIC,  ANTIMONY,  AND  BISMUTH 396 

XXIX.  SOME  IMPORTANT  ORGANIC  COMPOUNDS 421 

XXX.  SILICON  AND   BORON.     THE  ACID-FORMING  ELEMENTS  AND 

THE  PERIODIC  CLASSIFICATION : 428 

XXXI.  THE  PHYSICAL  PROPERTIES  OF  METALS.     ALLOYS 442 

XXXII.  THE  CHEMICAL  PROPERTIES  OF  METALS.     METALLURGY...  457 

XXXIII.  ELECTROCHEMISTRY 471 

XXXIV.  THE  PROPERTIES  OF  OXIDES,  HYDROXIDES,  AND  SALTS....  499 
XXXV.  SODIUM,  POTASSIUM,  RUBIDIUM,    AND    CAESIUM 515 

XXXVI.  CALCIUM,  STRONTIUM,  BARIUM.  AND  RADIUM 531 

vii 


viii  CONTENTS 

CHAPTER  PAGE 

XXXVII.  BERYLLIUM,  MAGNESIUM,  ZINC,  CADMIUM,  AND  MERCURY.  551 

XXXVIII.  ALUMINIUM 566 

XXXIX.  TIN  AND  LEAD 578 

XL.  COPPER,  SILVER,  AND  GOLD 588 

XLI.  IRON,  COBALT,  AND  NICKEL 603 

XLII.  THE  PLATINUM  METALS 623 

XLIII.  CHROMIUM,  MOLYBDENUM,  TUNGSTEN,  AND  URANIUM 626 

XLIV.  MANGANESE 639 

XLV.  RADIOACTIVITY.     THE  STRUCTURE  OF  ATOMS 646 

APPENDIX.  .             659 


INORGANIC  CHEMISTRY 


CHAPTER  I 
INTRODUCTION 

1.  What  is  chemistry,  and  why  should  the  science  be  studied 
by  one  who  is  seeking  a  liberal  education?  If  time  of  the  greatest 
value  is  to  be  given  to  the  study  of  -&  subject,  there  should  be  an 
appreciation  at  the  outset  of  what  is  to  be  gained  by  the  effort. 
Will  the  knowledge  acquired  and  the  logical  - inpde  of,  thinking 
developed  be  worth  the  labor?  Shall  we  know  the  things  about 
us  better  and  thus  increase  the  pleasure  of  living?  Let  us  first 
see  what  chemistry  is. 

Countless  changes  are  taking  place  in  the  material  world 
around  us.  In  many  of  these,  substances  are  undergoing  pro- 
found transformations.  Wood  and  coal  burn  and  are  changed 
into  ashes;  iron  rusts  and  falls  to  a  powder;  sweet  milk  becomes 
sour;  a  seed  sprouts,  and  feeding  on  the  substances  in  the  soil 
and  in  the  air  transforms  them  into  a  luxuriant  plant;  and  the 
plant  dies  and  becomes  but  a  part  of  the  soil  again.  The  chemist 
studies  such  changes  as  these  and  records  the  properties  of  the 
substances  which  undergo  transformation  and  the  properties  of 
the  substances  formed;  he  discovers  under  what  circumstances 
such  changes  take  place,  and,  with  this  knowledge,  is  able  to 
direct  many  of  them  at  will.  He  does  more;  he  formulates  laws 
which  express  the  general  behavior  of  substances,  and,  guided  by 
these  laws,  brings  about  countless  changes  which  never  took 
place  before. 

For  centuries  men  with  inquiring  minds  have  observed  nature 
closely,  and  have  recorded  and  brought  into  systematic  relation- 
ship many  transformations  in  the  substances  which  make  up  the 


2  INORGANIC  CHEMISTRY  FOR  COLLEGES 

material  world.  This  knowledge  forms  the  basis  for  the  facts  and 
laws  of  chemistry.  The  science  is  thus  a  vast  storehouse  of  infor- 
mation of  the  greatest  interest  and  value.  It  tells  us,  for  example, 
what  happens  when  silver  tarnishes,  when  plants  grow,  and  when 
bread  turns  sour.  But  more  than  this — it  teaches  us  how  to 
prevent  the  metal  from  tarnishing,  or  how  to  remove  the  stain 
once  it  is  formed;  it  shows  us  what  to  add  to  the  soil  to  facilitate 
the  growth  of  the  plant,  and  how  to  make  better  bread,  without 
the  risk  of  the  formation  of  the  undesired  acids. 

2.  The  study  of  the  changes  which  take  place  in  nature  of 
themselves,  is  but  a  small  part  of  the  science  of  chemistry.     New 
substances  of  the  greatest  value  have  been  made  by  the  chemist, 
and  ways  have  been  found  to  produce  new  and  desirable  effects. 
The  chemist  has  found  out  how  to  convert  iron  into  a  steel  which 
possesses  the  properties  required  in  tool  making;   he  has  made  a 
steel  for  armor  plate  designed  to  resist  most  effectively  the  impact 
of  projectiles;    he  has  produced  a  third  variety  with  properties 
which  adap'i*/  it  to  the  requirements  exacted  in  the  building  of 
light  but  strong  automobile  parts.     The  chemist  has  changed 
agriculture  from  an  art  to  a  science,  and  has  converted  disagree- 
able waste  products  into  valuable  plant  foods.     He  has  made 
available  processes  in  photography  which  place  it  among  the  fine 
arts;   he  has  furnished  antiseptics  which  prevent  decay  and  dis- 
ease.    The  chemist  builds  electric  batteries,  bleaches  cotton  and 
wool,  makes  cements  which  displace  wood  and  stone  as  building 
materials,  and  transforms  the  tar  obtained  by  heating  coal  into 
dyes  of  every  hue,  into  medicines,  and  into  explosives  of  the 
greatest  power.     We  truly  live  in  the  age  of  chemistry. 

Chemistry  plays  such  an  important  part  in  our  life  every  day, 
that  we  are  constantly  in  contact  with  it.  Life  itself  is  associated 
with  changes  in  the  materials  of  which  the  body  is  made  up. 
The  digestion  and  assimilation  of  foods  are  chemical  processes  of 
the  profoundest  interest.  The  clothes  we  wear,  and  nearly  every- 
thing which  adds  to  our  comfort  and  pleasure  in  life,  have  been 
under  the  hands  of  the  chemist  who  has  prepared  them  for  our 
use. 

3.  Will  not  a  knowledge  of  this  broad  and  wonderful  science 
add  to  our  mental  satisfaction  and  increase  the  pleasure  of  living? 
Have  we  not  advanced  in  our  education  when  we  have  learned  to 


INTRODUCTION  3 

interpret  changes  of  the  greatest  interest  which  are  taking  place 
all  around  us?  But  the  accumulation  of  useful  knowledge  is  not 
the  sole  benefit  gained  from  a  study  of  chemistry.  Other  results 
of  the  greatest  educational  value  follow,  when  the  study  of  the 
science  is  conducted  in  an  intelligent  manner.  In  the  laboratory 
the  student  examines  the  properties  of  typical  and  important 
substances,  and  investigates  the  transformations  of  these  sub- 
stances into  others  possessing  different  properties.  He  thus  has 
an  opportunity  to  observe  closely  and  to  record  his  observations. 
He  learns  to  see  things  as  they  are — a  power  which  comes  only 
with  training.  The  comprehension  and  the  use  of  the  laws  and 
theories  of  chemistry  develop  the  power  of  reasoning.  The  study 
of  cause  and  effect  and  the  relation  between  fact  and  theory, 
produce  logical  methods  of  thought  which  are  characteristic  of 
the  educated  man.  Careful  and  conscientious  work  on  the  part 
of  the  student  in  meeting  the  requirements  in  a  course  in  science 
will  yield  many  by-products  of  the  greatest  value.  He  will  develop 
promptness,  accuracy,  self-reliance,  and  clear  methods  of  thought 
and  expression.  In  the  case  of  students  who  do  not  continue  the 
study  of  chemistry  beyond  the  first  year,  these  by-products  prove 
of  greater  value  than  the  knowledge  of  facts  acquired.  Many 
of  the  latter  may  be  soon  forgotten,  but  the  mental  habits  formed 
persist. 

4.  Birth  of  Chemistry. — From  the  earliest  times  men  have 
observed  the  more  striking  changes  that  take  place  in  nature  and 
have  made  accidental  discoveries  of  great  value.  The  stone  age 
was  followed  by  the  age  of  bronze  when  the  way  to  work  metals 
was  found  out.  Many  of  the  substances  used  to-day,  which  are 
prepared  by  chemical  methods,  were  known  to  the  ancients,  but 
there  are  no  historical  records  of  their  discovery.  Glass,  for 
example,  which  has  been  known  for  over  3000  years,  was  prob- 
ably first  made  in  China.  According  to  tradition  a  method  of 
making  glass  was  discovered  independently  by  some  Phoenician 
sailors,  who  were  preparing  their  food  on  a  sandy  beach.  They 
supported  their  cooking  utensils  on  pieces  of  crude  soda  taken 
from  the  cargo  of  their  vessel.  The  heat  of  the  fire  brought  about 
a  reaction  between  the  sand  and  soda,  and  a  transparent  glassy 
substance  was  formed.  According  to  another  story  the  dis- 
covery was  made  by  Egyptian  priests  in  connection  with  the 


4  INORGANIC  CHEMISTRY  FOR  COLLEGES 

extraction  of  gold  from  the  earthy  material  with  which  it  was 
mixed.  They  added  soda  to  the  mixture  to  lower  the  tempera- 
ture to  which  the  mass  was  heated  to  melt  it,  and  found  that  a 
transparent  substance  was  formed.  Soap  was  also  made  in  the 
earliest  times;  wood  ashes  and  fat  were  used  and  it  is  easy  to  see 
how  the  process  could  be  discovered  by  accident.  Important 
discoveries  were  made  by  the  Egyptian  priests,  but  they  were 
kept  as  secrets,  and  many  of  them  were  forgotten  and  redis- 
covered later. 

But  advance  was  slow  as  long  as  observation  and  chance  alone 
led  the  way.  Directed  experimentation  was  necessary;  and  for 
this  a  motive  had  to  be  present.  It  was  with  the  rise  of  alchemy 
that  such  a  motive  appeared.  The  alchemists  sought  the  things 
which  were  supposed  to  lead  to  happiness — health  and  riches. 
They  endeavored  to  change  the  common  metals  into  gold,  hoping 
to  do  this  with  the  help  of  a  mysterious  substance,  called  the 
philosopher's  stone.  But  gold  without  health  is  of  little  value, 
so  a  search  was  made  for  the  elixir  of  life,  which  could  bring  back 
glorious  youth  to  the  aged.  Men  worked  at  these  problems  all 
over  Europe.  They  studied  everything  available,  mixed  things 
together,  and  heated  and  distilled  them  when  possible.  As  sub- 
stances appeared  to  affect  one  another  more  readily  when  they 
were  dissolved  in  some  liquid,  a  third  substance,  a  universal 
solvent,  was  sought  as  an  aid  toward  accomplishing  the  end  in 
view. 

Many  important  discoveries  were  made  as  the  result  of 
the  eager  search  by  the  alchemist  for  what  proved  to  be  unattain- 
able, and  some  of  the  processes  used  to-day  in  chemistry  were 
invented.  The  alchemists  refer  repeatedly  to  distillation,  extrac- 
tion, calcination,  coagulation,  etc.  They  prepared  and  studied 
many  of  the  compounds  which  are  to-day  commonly  used  in 
chemical  work.  Among  these  are  sulphuric,  nitric,  and  hydro- 
chloric acids,  alum,  soda,  ammonium  chloride,  niter,  and  com- 
pounds of  mercury,  arsenic,  and  antimony.  Some  of  the  alchem- 
ists gained  great  fame  as  a  result  of  their  work  and  writings. 
Geber,  an  Arab  who  lived  in  the  eighth  century,  exerted  an  influ- 
ence on  the  later  alchemists  of  the  Middle  Ages.  His  books  were 
translated  into  Latin  in  the  thirteenth  century  and  were  the 
standard  for  his  followers.  Other  noted  alchemists  were  Albertus 


INTRODUCTION  5 

Magnus  (1205-1280),  Roger  Bacon  (1214-1294),  and  Raymond 
Lullus  (1235-1315).  Alchemy  reached  its  height  in  the  thirteenth 
to  fifteenth  centuries.  It  was  customary  for  kings  and  princes 
to  have  alchemists  attached  to  their  courts,  for  the  rewards 
would  be  great  if  success  crowned  their  efforts.  There  was 
accumulated  in  this  way  a  vast  amount  of  information  upon 
which  the  science  of  chemistry  was  later  founded.  The  knowledge 
spread  slowly,  as  each  alchemist  guarded  his  secrets  carefully. 
In  order  to  gain  renown  for  his  accomplishments,  and  thus,  per- 
haps, win  a  place  at  court,  he  had  to  publish  to  the  world  his 
discoveries.  But  he  must  keep  his  secrets.  As  a  result,  alle- 
gorical and  figurative  language  was  used  in  such  a  way  that  both 
aims  were  reached;  the  discoveries  were  so  described  that  they 
appeared  of  the  greatest  importance,  but  at  the  same  time  the 
description  was  written  in  more  or  less  unintelligible  language. 
The  selfish  and  utilitarian  motive  which  guided  the  alchemist 
could  not  lead  to  the  development  of  a  science.  It  was  only 
when  a  desire  on  the  part  of  men  to  learn  the  truth  about  the 
wonders  of  nature  became  the  incentive,  that  chemistry  as  a 
science  was  born. 


CHAPTER  II 
PHYSICAL  AND  CHEMICAL  CHANGES 

5.  We  have  already  learned  that  chemistry  is  a  science  that 
deals  with  those  changes  in  nature  in  which  the  substances  before 
and  after  the  change  are  different.     When  iron  is  converted  into 
iron  rust  we  have  a  chemical  change;   for  the  powder  formed  is 
evidently  a  different  substance  from  the  hard  and  rigid  metal 
from  which  it  was  produced.     As  it  is  of  the  first  importance  to 
have  at  the  outset  as  clear  a  conception  as  possible  of  what  are 
classed  as  chemical  phenomena,  a  few  simple  experiments  will  be 
carefully  studied. 

6.  Physical    and    Chemical    Changes. — Let  us  consider  first 
what  happens  when  two  metals,  platinum  and  magnesium,  are 
heated  in  a  flame.     A  wire  of  platinum  when  heated  becomes  hot 
and  gives  off  light — we  could  see  it  in  a  dark  room.     A  change 
has  taken  place  in  it,  for  when  hot  it  is  red,  whereas  when  cold  it 
has  the  color  of  silver.     If  the  wire  is  removed  from  the  flame 
and  allowed  to  cool,  it  will  be  found  that  the  metal  has  all  the 
properties  it  possessed  before  being  heated — the  substance  has 
not  been  converted  into  a  new  substance  as  a  result  of  being 
heated.     If  we  carry  out  a  similar  experiment  with  magnesium, 
the  result  is  quite  different.     The  metal  burns  with  an  intense 
white  light,  and  is  changed  into  a  white  powder — a  new  substance 
is  formed.     This  experiment  with  the  two  metals  is  very  instruct- 
ive,  for  it  furnishes  examples  of  the  two  kinds  of  changes  which 
may  take  place  in  the  things  around  us.     In  the  case  of  the  plati- 
num the  change  did  not  yield  a  new  substance;  it  is  an  example 
of  what  is  called  a  physical  change.     In  the  case  of  the  mag- 
nesium, a  new  substance  was  formed  and  the  change  was  chemical. 
The  effect  of  heat  on  platinum  as  the  result  of  which  the  metal 
gives  off  light,  softens,  expands,  etc.,  is  studied  in  the  science  of 
physics.     The  change  of  metallic  magnesium,  when  heated,  into 

6 


PHYSICAL  AND  CHEMICAL  CHANGES  7 

the  white  powder,  which  is  called  magnesium  oxide,  is  studied  in 
the  science  of  chemistry. 

7.  In  the  experiment  just  described  the  results  were  brought 
about  by  the  application  of  heat.     Other  agencies  produce  physi- 
cal and  chemical  changes.     Two  experiments  will  illustrate  how 
so-called  mechanical  energy  can  produce  these  effects.     If  a  stick 
of  sulphur  is  brought  into  contact  with  a  pith-ball  suspended  by 
means  of  a  silken  thread,  the  substances  have  no  effect  on  each 
other.     If,  now,  the  sulphur  is  rubbed  briskly  with  a  piece  of 
silk  or  with  a  cat's  skin,  and  again  brought  into  contact  with  the 
ball,  it  will  be  found  that  the  two  substances  attract  each  other,  and 
the  ball  clings  to  the  sulphur.     If  the  two  are  now  separated  they 
repel  each  other;   when  the  sulphur  is  brought  near  the  pith-ball 
the  latter  moves  away  in  order  to  get  as  far  off  from  the  sulphur 
as  possible.     The  mechanical  energy  exerted  on  the  sulphur  has 
changed  it;   electricity  has  been  produced  on  its  surface,  and  this 
brings  about  the  difference  in  its  action  on  the  pith-ball.     But 
it  is  still  sulphur.     The  change  produced  is  a  physical  one  and  is 
studied  in  detail  in  a  course  in  physics.  .   . 

When  many  substances  are  rubbed  together  more  profound 
changes  take  place  than  that  which  was  the  result  of  the  experi- 
ment just  described.  Mercury  and  iodine  furnish  an  excellent 
example.  Mercury  is  a  heavy  liquid,  with  the  color  of  silver, 
and  is  recognized  as  the  substance  used  in  thermometers.  Iodine 
is  a  black  solid,  which  dissolves  in  alcohol,  forming  a  brown  liquid, 
known  as  tincture  of  iodine.  When  the  two  substances  are  rubbed 
together,  in  the  presence  of  a  little  alcohol,  which  makes  the 
change  take  place  more  rapidly,  they  are  transformed  into  a  red 
powder.  The  mercury  and  the  iodine  disappear  and  a  substance 
with  new  properties,  called  mercuric  iodide,  is  produced.  The 
mechanical  energy  expended  served  to  bring  the  two  substances 
into  intimate  contact,  and  a  chemical  change  took  place. 

8.  Light  is  a  form  of  energy  which  produces  changes.     Certain 
dyes  fade  in  sunlight;  the  art  of  photography  is  based  upon  the 
action  of  light  on  substances  containing  silver  and  other  metals; 
and  the  growth  of  plants  requires  this    form  of    energy.     Two 
simple  experiments  will  be  instructive.     A   radiometer,    Fig.    1, 
consists  of  two  rectangular  pieces  of  thin  sheet  metal  set  at  right 
angles  and  suspended  at  the  point  of  intersection  in  such  a  way 


8  INORGANIC  CHEMISTRY  FOR  COLLEGES 

that  the  whole  can  rotate  freely.  Alternate  sides  of  the  plates 
are  blackened.  This  arrangement  is  put  into  a  glass  vessel  from 
which  the  air  is  exhausted,  in  order  to  reduce  to  a  minimum  the 
friction  produced  when  the  plates  rotate.  When  a  radiometer 
is  placed  in  sunlight  the  vanes  rotate  rapidly.  The  light  falling 
on  the  blackened  surfaces  repels  them  and  thus  produces  the 
rotation.  The  effect  is  produced  by  radient  energy.  If  the 
heat  radiations  are  absorbed  by  passing  the 
light  through  a  proper  screen,  the  rotations 
continue.  Light  causes  in  this  case  a  phys- 
ical change. 

If  a  piece  of  paper  is  dipped  into  a  solution 
made  by  dissolving  silver  nitrate  in  water,  and 
is  exposed  to  light,  it  will  turn  black.  Silver 
nitrate  is  made  by  heating  metallic  silver 
with  nitric  acid;  it  is  a  white  crystalline 
substance.  When  exposed  to  sunlight,  in 
contact  with  paper,  it  is  decomposed  and  the 
metal  is  produced  in  a  form  which  is  black.  A 
new  substance  is  the  result  of  the  change; 
silver  nitrate  is  converted  into  silver.  Light 
in  this  case  effects  a  chemical  change. 

9.  Electricity  produces  physical  and  chemical  changes.  If  a 
wire  through  which  a  current  of  electricity  is  passing  is  brought 
over  a  compass  needle,  the  latter  will  be  deflected  and  will  no  longer 
point  to  the  north.  When  the  wire  is  removed  the  needle  returns 
to  its  original  position.  The  electricity  produced,  evidently,  a 
change  which  was  physical.  This  phenomenon  is  of  the  greatest 
importance,  and  has  been  studied  in  detail;  upon  it  is  based  the 
transformation  of  electricity  into  mechanical  power  as  exemplified 
in  the  electric  motor. 

A  chemical  change  can  be  readily  produced  by  means  of  elec- 
tricity. If  the  ends  of  two  platinum  wires  connected  with  a 
source  of  electricity  are  placed  in  water  which  contains  a  little 
acid,  bubbles  of  gas  will  be  seen  to  form  on  the  wires  in  the  solu- 
tion, and  to  rise  to  the  surface.  Water  does  not  allow  electricity 
to  flow  through  it  readily  and  the  small  amount  of  acid  is  added 
te  make  it  possible  for  the  current  to  flow.  If  the  electricity  is 
allowed  to  pass  through  the  solution  for  a  long  enough  time,  the 


PHYSICAL  AND  CHEMICAL  CHANGES 


9 


water  will  gradually  disappear  and  large  amounts  of  the  gases 
will  be  produced.  The  change  produced  is,  evidently,  a  chemical 
one;  the  water  is  converted  into  other  substances,  which  possess 
entirely  different  properties  from  it.  A  piece  of  apparatus  which 
is  commonly  used  to  demonstrate  this  experiment  is  shown  in 
Fig.  2.  It  was  devised  by  the  chemist  Hofmann  and  is  known 
by  his  name.  The  ends  of  the  wire  connected  with  the  source  of 
electricity  are  joined  to  the  apparatus  at  A  and  B.  At  these  points 
wires  pass  through  the  glass  and  terminate 
in  platinum  plates,  which  are  called  the 
electrodes.  The  gas  that  rises  from  the 
electrodes  collects  in  the  two  tubes  and  forces 
the  water  up  into  the  receiver  C.  The  gases 
formed  can  be  removed  from  the  apparatus 
and  collected  by  opening  the  stop-cocks  at  the 
ends  of  the  tubes.  The  Hofmann  apparatus 
is  a  very  convenient  one  with  which  to  study 
decompositions  yielding  gases  produced  by 
passing  electricity  through  solutions.  We  shall 
have  occasion  to  refer  to  it  repeatedly. 

Of  the  four  kinds  of  energy  which  bring 
about  physical  and  chemical  change,  heat  and 
electricity  are  used  extensively  in  chemistry. 
The  application  of  light  is  limited  at  present 
largely  to  photographic  processes.  Mechani- 
cal energy  is  used  to  some  extent  when  chemical 
changes  between  gases  are  brought  about, 
for  when  the  latter  are  highly  compressed  the  results  are,  in  cer- 
tain cases,  more  satisfactory  than  under  ordinary  conditions. 

10.  Variable  and  Characteristic  Properties. — There  is  much 
that  can  be  learned  from  a  consideration  of  the  experiments  just 
described.  We  have  seen  that  the  changes  brought  about  in 
matter  can  be  divided  into  two  classes:  those  in  which  a  sub- 
stance is  converted  into  another  substance — a  chemical  change, — 
and  those  in  which  no  new  substance  appears — a  physical  change. 

In  order  to  determine  to  which  class  any  given  change  belongs, 
it  is  necessary  to  note  carefully  the  properties  of  all  materials 
involved  before  and  after  the  change.  We  recognize  the  different 
forms  of  matter  by  their  properties.  The  common  ones  made 


FIG.  2. 


10  INORGANIC  CHEMISTRY  FOR  COLLEGES 

use  of  are  those  which  appeal  directly  to  our  senses.  Through 
this  means  we  recognize  form,  luster,  solubility,  and  color;  hard- 
ness, brittleness,  and  ductility;  odor  and  taste.  Examine  a  crys- 
tal of  sugar — a  piece  of  rock  candy — and  see  how  it  can  be  identi- 
fied by  its  properties.  It  has  a  definite  shape  and  shiny  surfaces, 
is  transparent,  colorless,  and  soluble  in  water,  is  hard  and  brittle, 
is  odorless,  and  possesses  a  sweet  taste.  The  identity  of  this 
particular  piece  of  sugar  is  determined  by  all  these  properties; 
we  could  recognize  it  among  other  things  and  other  pieces  of  sugar. 
But  must  all  samples  of  sugar  possess  all  of  these  properties? 

Suppose  we  grind  the  crystal  to  a  fine  powder.  The  form  has 
changed,  it  is  not  transparent,  and  we  cannot  see  shiny  surfaces; 
it  no  longer  possesses  many  properties  of  the  crystal.  Is  it  still 
sugar?  Has  a  new  substance  been  formed?  Was  the  change  a 
physical  or  chemical  one?  The  powder  has  two  properties  which 
the  crystal  possessed;  it  dissolves  in  water  and  the  solution  has 
a  sweet  taste.  We  know  it  is  sugar;  and  the  change  is  a  physical 
one.  It  is  evident,  therefore,  that  we  must  differentiate  two  kinds 
of  properties — those  which  serve  to  identify  any  particular  piece 
of  matter,  which  is  called  a  body,  and  those  which  are  associated 
with  the  material  of  which  the  body  is  made  up.  Properties 
of  the  former  class  are  called  variable  properties,  and  those  of 
the  latter,  characteristic  The  shape  of  a  piece  of  sugar  is  a  vari- 
able property;  its  sweet  taste  when  dissolved  in  water  is  a 
characteristic  property.  When  chemical  changes  take  place,  sub- 
stances are  formed  which  possess  different  characteristic  properties 
from  those  of  the  substances  entering  into  the  change.  A  difference 
in  a  single  property  of  this  type  shows  that  another  substance 
has  been  formed. 

In  enumerating  characteristic  properties  it  is  necessary  to 
state  the  external  conditions  to  which  the  substance  is  subjected. 
A  characteristic  property  of  water  is  that  it  is  a  liquid;  but  this 
statement  is  found  to  be  true  only  when  water  is  examined  within 
a  certain  range  of  temperature;  below  the  freezing-point  it  becomes 
a  solid.  Referring  back  to  our  experiment  with  the  platinum 
wire,  it  will  be  remembered  that  the  metal  possessed  the  color  of 
silver  before  being  heated,  and  that  when  the  change  was  brought 
about  this  disappeared  and  the  wire  became  red.  Was  the  change 
physical  or  chemical?  The  most  evident  change  in  property  was 


PHYSICAL  AND  CHEMICAL  CHANGES  11 

in  color.  We  must  compare  properties  under  the  same  external 
conditions;  so  we  let  the  wire  cool  and  examine  it.  It  has  the 
color  of  silver  and  has  the  other  characteristic  properties  it  pos- 
sessed before  being  heated;  the  change  is  thus  a  physical  one. 

11.  Physical     Properties     and    Chemical    Properties. — The 
properties  which  substances  possess  are  classified,  at  times,  from 
a  point  of  view  different  from  that  used  in  dividing  them  into 
accidental  and  characteristic.     A  distinction  is  drawn  in  this  case 
between  the  properties  which  we  recognize  directly,  or  when  a 
physical  change  takes  place,  and  those  which  become  evident 
only  when  a  chemical  change  occurs;  the  former  are  called  physi- 
cal properties,  the  latter  chemical  properties.     Some  examples 
will  make  this  distinction  clear.     Iron  is  hard,  can  be  drawn  out 
into  a  wire,  is  heavy,  and  allows  electricity  to  pass  through  it;  it 
has,  thus,  the  properties  of  hardness,  ductility,  density,  and  elec- 
trical conductivity.     All  these  properties  can  be  recognized  with- 
out a  chemical  change  taking  place  in  the  metal;  they  are  physical 
properties.     Iron  rusts  when  exposed  to  the  air,  and  dissolves 
when   put   into  an  acid,  such  as  vinegar,  for  example.     These 
properties  become  evident  only  when  the  iron  undergoes  a  chem- 
ical change;    they  are  chemical  properties.     Since  we  recognize 
substances  by  both  physical  and  chemical  properties,  attention 
must  be  paid  to  the  two  classes  in  the  study  of  chemistry. 

The  distinction  which  has  just  been  drawn  is  often  expressed 
in  a  different  way.  Physical  properties  are  classed  simply  as 
properties,  and  what  we  have  called  the  chemical  properties  of  a 
substance  are  -referred  to  as  its  chemical  conduct.  The  so-called 
chemical  properties,  as  has  been  said,  become  evident  only  when 
the  substance  possessing  them  enters  into  chemical  reaction; 
they  are  an  expression  of  its  chemical  behavior  or  conduct  when 
undergoing  chemical  change.  This  method  of  expressing  the 
classification  is  perhaps  the  better  one. 

12.  Matter  and  Energy. — A  great  deal  can  be  learned  from 
the  simple  experiments  which  were  described  earlier  in  the  chapter. 
It  will  be  recalled  that  the  changes  were  brought  about  by  means 
of  heat,   mechanical  energy,   light,   and  electricity.    These  are 
what  are  known  as  forms  of  energy.     It  is  necessary  to  differ- 
entiate between  matter  and  energy.     A  strictly  accurate  defini- 
tion of  matter  is  difficult  to  formulate.     Our  experience  and  com- 


12  INORGANIC  CHEMISTRY  FOR  COLLEGES 

mon  sense  furnish  us  with  a  conception  of  matter.  Matter 
occupies  space,  it  has  inertia,  that  is,  it  requires  force  to  move  it; 
it  is  the  stuff  of  which  the  universe  is  made.  Energy,  on  the  other 
hand,  is  non-material;  we  become  conscious  of  it  only  when  it  is 
associated  with  matter.  A  stone  held  in  the  air  is  different  from 
the  same  stone  resting  on  the  earth;  for  by  allowing  the  former 
to  drop  we  can  obtain  work  from  it,  drive  a  nail,  or  crush  grain. 
The  stone  held  away  from  the  earth  is  said  to  have  potential 
energy;  it  gives  it  up  when  it  falls;  and  to  raise  it  from  the  earth 
back  to  its  original  position,  work  must  be  done  upon  it.  Energy 
manifests  itself  in  work. 

In  our  experiments  we  have  made  use  of  four  important  forms 
of  energy:  heat,  light,  electricity,  and  mechanical  energy.  These 
can  all  be  made  to  do  work.  The  various  forms  of  energy  can  be 
transformed,  one  into  the  other.  For  example,  the  mechanical 
energy  of  a  water-fall  can  be  transformed  into  electricity  by  means 
of  a  water-wheel  and  a  dynamo.  This  can,  in  turn,  be  converted 
into  light,  heat,  or  mechanical  energy. 

13.  Matter  and  energy  are  always  associated.  When  any 
change  occurs  there  is  always  a  change  in  the  energy;  there  may 
or  may  not  be  a  change  in  the  matter.  From  this  point  of  view 
we  can  define  physical  and  chemical  change;  if  the  change  con- 
sists solely  in  energy  it  is  physical,  if  the  matter  changes  it  is  chemi- 
cal. Physics  is  thus  the  science  which  has  to  do  with  the  study 
of  changes  in  energy,  either  in  amount  or  kind.  As  energy  becomes 
manifest  only  through  its  action  on  matter,  the  physicist  studies 
the  behavior  of  substances  which  do  not  alter  hi  composition 
when  energy  changes  take  place  in  them.  Physicists  have  studied, 
for  example,  the  effect  produced  by  moving  a  wire  before  a  magnet, 
and  have,  as  a  consequence,  invented  the  dynamo.  They  have 
investigated  the  behavior  of  light  when  it  passes  through  glass, 
and  the  telescope  is  the  result. 

In  chemical  changes  both  matter  and  energy  change.  The 
older,  descriptive  chemistry  busied  itself  with  the  matter  involved 
only.  In  modern  chemistry,  the  energy  changes  which  occur 
simultaneously  with  the  changes  in  matter  have  been  the  subject 
of  much  study.  This  branch  of  the  science  is  aptly  called  physical 
chemistry.  For  convenience  it  has  been  divided  into  several 


PHYSICAL  AND  CHEMICAL  CHANGES  13 

divisions:  thermochemistry  treats  of  the  changes  in  heat  energy 
when  matter  changes;  photochemistry  has  to  do  with  the  rela- 
tionship between  light  and  matter;  in  electrochemistry  is  studied 
the  effect  of  electricity  in  producing  chemical  changes,  as  well  as 
the  production  of  electricity  through  chemical  means;  and  other 
branches  of  the  science  consider  the  effect  produced  by  mechanical 
energy. 

The  subject  of  chemistry  is  such  a  broad  one  that  it  has  been 
found  desirable  to  study  it  from  many  points  of  view.  Inorganic 
chemistry  has  to  do  with  the  substances  found  in  the  so-called 
mineral  kingdom.  The  chief  interest  in  organic  chemistry  centers 
in  the  compounds  formed  as  the  result  of  life-processes,  and  in 
the  substances  prepared  in  the  laboratory  from  these  compounds. 
Physiological  chemistry,  as  the  name  implies,  is  the  chemistry  of 
the  processes  studied  in  physiology.  Many  other  branches  of 
the  science  are  highly  developed,  such  as  industrial,  metallurgical, 
mineralogical  chemistry,  etc. 


EXERCISES 

1.  State  as  fully  as  possible  the  accidental  and  characteristic  properties 
of  the  following:    (a)  a  crystal  of  salt,  (6)  a  copper  wire,  and  (c)  a  block  of 
wood. 

2.  Study  the  experiments  described  to  illustrate  physical  and  chemical 
change  (sections  6-9),  and  state  the  accidental  and  characteristic  properties 
of  the  substances  used  and  those  obtained  in  each  case.     Show  why  the  con- 
clusions reached  are  justified.       • 

3.  State  which  of  the  following  are  physical  changes  and  which  are  chemi- 
cal changes.     Examine  each  case  from  the  point  of  view  of  characteristic 
properties,  (a)  tearing  a  piece  of  paper  into  small  bits,  (6)  the  falling  of  a  stone, 
(c)  burning  of  wood,    (d)  digestion  of  food,    (e)  lighting  an   incandescent 
lamp,  (/)  collision  of  two  railroad  trains,  (gr)  change  of  cider  into  vinegar, 
(h)  making  toast  from  bread,  (i)  exploding  fire-works,  (j)  decay  of  a  flower. 

4.  Name  a  physical  and  chemical  change  which  occurs,  (a)  when  a  motor 
boat  with  a  gasoline  engine  is  running,  (6)  when  a  bell  is  rung  by  an  electric 
battery,  (c)  when  a  lamp  burns. 

5.  Name  three  physical  and  three  chemical  changes  not  mentioned  in 
this  book  and  give  reasons  for  your  conclusion  in  each  case. 

6.  Name  as  fully  as  possible  the  physical  properties  of,  (a)  starch,  (6)  salt, 
(c)  gold,  (d)  lead,  and  the  chemical  properties  of  (e)  paper,  (/)  copper,  (g)  iron. 

7.  How  could  the  following  changes  in  energy  be  brought  about:   (a)  heat 


14  INORGANIC  CHEMISTRY  FOR  COLLEGES 

into  light,    (6)   heat  into  mechanical  energy,    (c)   electricity  into  light,    (d) 
mechanical  energy  into  heat,  (e)  electricity  into  heat? 

8.  Name  one  change  in  each  case  not  mentioned  in  the  text  which  can  be 
produced  by  (a)  heat,  (6)  mechanical  energy,  (c)  light,  (d)  electricity.     State 
which  of  the  changes  are  physical  and  which  are  chemical. 

9.  State  what  changes  in  energy  take  place  in  each  of  the  following: 
(a)  driving  a  nail  with  a  hammer,  (6)  running  an  electric  car  uphill,  (c)  heating 
an  electric  flat  iron. 


CHAPTER  III 
ELEMENTS  AND  COMPOUNDS 

14.  As  the  result  of  the  study  of  the  last  chapter  the  student 
has  begun  to  learn  how  to  observe  and  analyze  phenomena.     A 
closer  inspection  and  a  deeper  understanding  of  the  nature  of  the 
substances  involved  are  necessary.     Again,  we  shall  seek  the  help 
of  experiments  in  order  to  make  clear  the  distinctions  to  be  drawn. 
We  shall  repeatedly  do  this,  for  experimentation  is  the  foundation 
upon  which  chemistry  is  based.     We  solve  our  problems,  and 
clarify  our  conceptions  by  direct  questions  to  nature;   we  bring 
together   the   things   whose   behavior   we   wish   to   understand, 
observe  for  ourselves  the  result,  and  draw  our  own  conclusions. 
This  is  the  so-called  scientific  method;   it  is  a  more  or  less  novel 
one  to  the  student  whose  point  of  view  has  been  formed  largely 
by  reliance  on  authority.     He  has  drawn  his  knowledge  from 
books,  and  has  had  little  opportunity  to  find  out  for  himself. 
In  chemistry  he  will  learn  to  see  with  his  own  eyes;   to  answer 
many  of  his  questions  by  obtaining  first-hand  knowledge  in  the 
laboratory  and  through  the  experimental  demonstrations  in  the 
lecture-room.     It  should  be  constantly  borne  in  mind  that  chem- 
istry is  correctly  studied  in  this  way,  and  the  student  will  be 
taught  to  realize  the  great  value  of  such  a  method  in  all  problems 
he  may  meet  inside  or  outside  the  classroom. 

15.  Pure    Substances    and    Mixtures. — When    we    examine 
closely  a  piece  of  marble  and  a  piece  of  granite  we  observe  a  marked 
difference  between  them.     The  one  appears  to  be  uniform  in 
properties,   the  other  not.     We  can  distinguish  in  the  granite 
three  distinct  substances:   one  is  white,  another  is  pink,  and  the 
third  is  black  and  glistens.     Our  common  sense  tells  us  that 
granite  is  a  mixture,  whereas  marble  appears  to  be  a  uniform 
substance.     If  we  were  to  powder  the  two  specimens  we  could 
separate  the  three  constituents  of  the  granite;  in  the  case  of  the 

15 


16  INORGANIC  CHEMISTRY  FOR  COLLEGES 

marble,  however,  each  piece  would  be  made  up  of  the  same  material 
as  every  other  piece.  Granite  is  a  mixture,  and  marble  is  what 
is  called  in  chemistry  a  pure  substance.  In  a  pure  substance 
every  part  is  like  every  other  part;  the  substance  is  uniform  in 
properties.  This  refers,  of  course,  to  what  we  have  called  char- 
acteristic properties.  Two  pieces  of  marble  might  possess  differ- 
ent sizes  and  shapes,  and  yet  they  are  both  marble;  their  character- 
istic properties  are  identical. 

We  have  chosen  a  very  simple  example  to  bring  out  the  mean- 
ing of  the  words  mixture  and  pure  substance.  We  made  use  of 
the  characteristic  property  of  color  in  this  case  to  come  to  the 
conclusion  that  granite  is  a  mixture.  If  we  were  asked  to  deter- 
mine whether  a  given  material  uniform  in  color,  is  a  pure  sub- 
stance or  a  mixture,  what  would  we  do  to  aid  us  in  answering  the 
question?  We  must,  evidently,  make  use  of  a  characteristic 
property,  other  than  that  of  color.  If  we  had  a  mixture  made  up 
of  sand  and  sugar,  both  in  the  form  of  a  very  fine  powder,  we  could 
readily  show  a  lack  of  uniformity  by  using  solubility  in  water  as 
the  characteristic  property.  The  sugar  would  dissolve,  and  the 
sand  would  not;  the  material  is  a  mixture,  for  different  parts  of 
it  possess  different  characteristic  properties. 

16.  Another  experiment  clearly  illustrates  the  point  under 
discussion.  Let  us  mix  some  iron  and  some  sulphur,  both  finely 
powdered,  and  grind  them  together.  We  obtain  as  the  result, 
what  appears  to  the  casual  observer  to  be  a  uniform  product.  Is  it 
a  mixture  or  a  pure  substance?  In  this  case  we  can  make  use  of 
one  of  a  number  of  characteristic  properties  to  answer  the  question. 
Iron  is  gray,  and  sulphur  is  yellow.  With  a  magnifying  glass  we 
can  observe  gray  and  yellow  particles;  the  product  is  a  mixture. 
Again,  iron  is  attracted  by  a  magnet,  and  sulphur  is  not.  If  we 
bring  a  magnet  near  the  product,  a  part  only  clings  to  it.  A 
third  method  can  be  used.  Sulphur  dissolves  in  a  liquid  which 
is  called  carbon  disulphide,  while  iron  does  not  dissolve.  If  we 
shake  the  product  with  this  liquid  and  then  pour  the  mixture  on 
a  piece  of  paper  which  is  folded  and  placed  in  a  funnel,  a  clear 
liquid  will  run  through  the  paper  and  the  iron  will  not.  When 
this  liquid  is  allowed  to  evaporate,  it  will  be  found  that  sulphur 
is  left  behind.  We  have  by  this  means  separated  our  product 
into  two  distinct  parts,  each  of  which  is  different  from  the  other. 


ELEMENTS  AND  COMPOUNDS  17 

17.  Chemical  Reactions. — By  carrying  our  experiment  with 
iron  and  sulphur  further,  we  can  learn  much  more.     A  certain 
weight  of  sulphur  and,  of  iron,  which  experience  has  shown  are 
the  correct  amounts,  are  mixed  and  heated.     The  tube  contain- 
ing the  mixture  is  held  in  such  a  position  that  the  bottom  end  of 
it,  only,  is  placed  in  the  flame  of  the  burner.     In  a  short  time  the 
mixture  will  get  red  hot  where  it  is  being  heated.    The  tube  is 
then  removed  from  the  flame.     A  striking  phenomenon  takes 
place;  the  brilliant  glow  slowly  travels  up  the  tube,  and,  finally, 
the  whole  mixture  gives  off  light.     When  the  tube  is  cold  it  is 
broken  and  the  product  is  examined.     Is  it  still  a  mixture  of  iron 
and  sulphur?    We  can  apply  the  tests  used  before.     The  prod- 
uct now  is  no  longer  a  powder.     In  the  lump  which  has  been 
formed  we  cannot  distinguish  the  fine  particles  of  sulphur;    a 
magnet  has  no  effect  on  it;   carbon  disulphide  does  not  dissolve 
from  it  yellow  sulphur.     A  profound  change  has  occurred;    the 
product  lacks  the  characteristic  properties  of  iron,  and  of  sulphur. 
The  change  has  been  a  chemical  one — there  has  been  a  change  in 
substance,  and  energy  has  been  transformed;   for  as  a  result  of 
the  change  light  and  heat  were  produced.     A  chemical  reaction 
has  occurred.     What  took  place  is  called  a  reaction  because  the 
iron  acted  on  the  sulphur,  and  the  sulphur  acted  on  the  iron. 

18.  Chemical  Compounds. — We  next  ask  what  has  become 
of  the  iron  and  of  the  sulphur.     The  two  substances  have  dis- 
appeared as  such,  and  a  new  substance  has  been  formed.     The 
transformation  has  been  most  remarkable,  for  we  can  no  longer 
recognize  iron  or  sulphur  in  the  product  of  the  reaction.     We 
explain  this  by  saying  the  two  substances  have  united  chemically; 
a  chemical  compound  has  been  formed.     There  was  a  marked 
change  in  the  energy  as  well  as  in  the  matter  when  the  reaction  took 
place.     In  order  to  bring  about  the  change  quickly  it  was  neces- 
sary only  to  start  the  reaction;   it  then  proceeded  of  itself,  and 
light  and  a  large  amount  of  heat  were  produced.     Whence  came 
this  energy?    Experience,  based  on  more  than  a  century  of  care- 
ful experimentation,  teaches  us  that  we  cannot  make  energy; 
the  amount  in  the  universe  is  constant,  and  man  can  only  change 
it  from  one  form  to  another.     As  this  is  a  fundamental  truth, 
the  energy  transformed  must  have  been  associated  with  the  iron 
and  the  sulphur.     Energy  in  this  form  can  be  recognized  only 


18  INORGANIC  CHEMISTRY  FOR  COLLEGES 

when  the  substances  containing  it  undergo  chemical  change;  it 
is  then  transformed,  in  part,  into  other  forms  of  energy.  We  are 
thus  led  to  add  a  new  form  of  energy  to  those  about  which  we  have 
studied;  it  is  called  chemical  energy. 

In  the  light  of  these  conceptions  we  can  describe  more  fully 
what  occurs  when  iron  reacts  with  sulphur.  Iron  as  we  know  it 
consists  of  a  certain  form  of  matter  with  which  is  associated 
chemical  energy;  likewise  sulphur,  which  possesses  the  properties 
we  recognize,  consists  of  a  different  form  of  matter  associated 
with  energy.  When  the  two  unite  chemically  a  part  of  the  energy 
is  lost;  it  is  transformed  into  light  and  heat.  The  resulting 
compound  contains  the  material  substance  of  the  iron  and  that 
of  the  sulphur,  but  less  energy  than  the  two  possessed  before  the 
change.  The  properties  of  substances  change  when  we  change 
the  energy  they  contain,  although  the  amount  of  material  of  which 
the  substance  is  made  does  not  alter. 

The  compound  formed  when  iron  reacts  with  sulphur  is  called 
iron  sulphide;  the  name  is  well  chosen  as  it  tells  us  that  iron  and 
sulphur  are  present  in  it.  This  compound  is  a  pure  substance; 
we  can  study  it  in  many  ways  and  we  shall  find  that  every  part 
of  it  has  the  same  characteristic  properties  possessed  by  every 
other  part. 

When  iron  and  sulphur  are  brought  together  we  have  seen  that 
a  mixture  is  formed;  when  this  is  heated  a  chemical  reaction 
takes  place,  chemical  energy  is  lost,  being  transformed  into  light 
and  heat,  and  a  chemical  compound  results. 

19.  One  more  interesting  question  remains  to  be  answered. 
Can  we  separate  the  compound  into  its  constituents  by  any  means? 
Can  we  get  back  the  iron  and  the  sulphur?  We  shall  see  later 
that  this  can  be  done,  but  it  is  of  importance  to  note  that  in  order 
to  do  this  we  must  make  use  of  chemical  reactions;  we  must  give 
back  to  the  iron  and  to  the  sulphur  the  energy  which  they  lost 
when  they  united  with  each  other.  If  we  restore  to  the  matter 
which  is  present  in  iron  the  right  amount  of  chemical  energy,  it 
becomes  metallic  iron  again.  We  thus  see  that  it  is  possible  to 
separate  both  mixtures  and  compounds  into  their  constituents. 
In  the  case  of  mixtures  we  use  mechanical  or  physical  means;  in 
the  case  of  compounds  we  must  resort  to  chemical  changes.  For 
this  reason  the  two  classes  have  been  called  mechanical  mixtures 


ELEMENTS  AND  COMPOUNDS  19 

and  chemical  compounds.  We  are  again  impressed  with  the 
necessity  of  a  clear  understanding  of  what  is  meant  by  a  physical 
change  and  a  chemical  change.  The  discussion  has  been  gone 
into  at  such  length  on  account  of  its  fundamental  importance. 

20.  Elements. — There  is  still  more  we  can  learn  from  the 
experiment  with  iron  and  sulphur.  The  substance  formed  as  the 
result  of  their  interaction  contains  both  sulphur  and  iron;  two 
distinct  forms  of  matter  enter  into  its  composition.  But  what  of 
the  iron  itself?  Does  it  contain  more  than  one  substance?  For 
many  years  chemists  have  studied  iron  and  they  have  never 
obtained  from  it  alone,  more  than  one  kind  of  matter.  It  is, 
therefore,  called  an  element.  Sulphur  is  also  an  element;  it  has 
never  been  broken  up  into  two  or  more  substances. 

Iron  sulphide,  on  the  other  hand,  is  a  compound,  for  by  using 
the  right  methods  it  can  be  decomposed  into  iron  and  sulphur, 
the  elements  of  which  it  is  made  up.  A  chemical  compound  has 
the  characteristics  of  a  simple  substance,  that  is,  it  is  uniform  in 
properties;  in  addition,  it  contains  two  or  more  elements  in  chemical 
combination.  It  is  important  to  note  that  the  meaning  of  the 
word  compound  as  used  in  chemistry,  is  quite  different  from  that 
commonly  assigned  to  it — to  the  ordinary  definition  as  given  in 
the  dictionary. 

Centuries  of  study  have  led  to  the  conclusion  that  the  universe 
is  made  up  of  about  ninety  elements.  The  infinite  number 
of  substances  in  the  world  have  been  produced  as  the  result  of  the 
combination  in  different  proportions  of  these  elements.  Only  a 
small  number  of  the  elements  are  present  in  the  things  with  which 
we  come  in  contact.  Less  than  a  dozen  are  contained  in  the  great 
variety  of  living  matter  around  us. 

EXERCISES 

1.  State  which  of  the  following  are  pure  substances  and  which  mixtures: 
(a)  salt,  (6)  vinegar,  (c)  sand,  (d)  sugar,  (e)  paper,  (/)  baking  soda,  (g)  bread, 
(A)  an  egg. 

2.  How  could  you  show  that  the  following  were  mixtures  and  how  would 
you  separate  each  into  its  constituents?     (a)  chalk  and  salt,  (6)  powdered 
iron  and  powdered  coal,  (c)  a  shoe  blacking  made  of  soot  and  grease. 

3.  Name  three  substances  that  are  elements  and  three  that  are  compounds. 

4.  Write  out  definitions  of  the  following  using  your  own  words :   (a)  physi- 


20  INORGANIC  CHEMISTRY  FOR  COLLEGES 

cal  change,  (6)  chemical  change,  (c)  mixture,  (d}  element,  (e)  chemical  com- 
pound, (/)  chemical  energy. 

5.  State  what  changes  in  energy  take  place  when  (a)  wood  burns,  (6) 
water  falls,  (c)  electricity  is  made  by  a  dynamo,  (d)  electricity  is  made  by  a 
dry  cell,  (e)  an  electric  car  moves,  (/)  toast  is  made  on  an  electric  stove,  (g) 
magnesium  is  used  in  a  flash-powder  in  photography  at  night,  and  (A)  when 
gasoline  is  exploded  in  a  motor. 


CHAPTER  IV 
OXYGEN 

21.  Fire  has  been  a  source  of  fear,  wonder,  and  veneration 
from  prehistoric  times;  it  has  aroused  the  keenest  interest  and 
curiosity  in  man,  and  has  been  the  subject  of  speculation  by  phi- 
losophers. It  has  received  the  closest  study  by  investigators, 
and,  as  a  result,  its  secret  has  been  found  out,  and  we  have  learned 
how  to  master  it.  The  Greeks  were  of  the  opinion  that  the  uni- 
verse was  made  up  of  four  principles:  earth,  air,  fire,  and  water. 
But  the  philosophy  of  the  ancients  did  not  get  them  very  far 
toward  the  solution  of  the  secrets  of  nature.  They  wasted  time 
trying  to  solve  the  mysteries  of  natural  processes  by  pure  reason- 
ing alone.  They  disputed  at  great  length,  for  example,  as  to 
whether  matter  was  infinitely  divisible,  a  purely  philosophical 
question,  but  did  not  try  to  find  out  more  about  matter  itself  by 
examining  it.  One  philosopher  is  said  to  have  had  his  eyes  put 
out,  so  that  he  might  reason  the  more  keenly  and  not  be  distracted 
by  the  crude  material  world.  Such  methods  were  futile,  and  after 
centuries  of  ignorance  man  learned  the  art  of  experimentation. 
Then  the  truth  was  discovered,  the  mystery  disappeared,  and 
fire  ceased  to  be  worshiped;  it  became  the  servant  of  man. 

If  we  are  to  produce  fire  by  burning  wood,  we  know  air  must 
be  present.  It  was  the  discovery  and  isolation  of  the  substance 
in  the  air  which  made  burning  possible,  that  solved  the  mystery 
of  fire.  Historical  research  has  disclosed  the  fact  that  a  number 
of  chemists  discovered  at  about  the  same  time  this  important 
constituent  of  the  air — the  substance  which  was  later  called  oxygen. 
As  early  as  1489  the  observation  was  made  that  when  the  sub- 
stance we  now  call  the  oxide  of  mercury  is  heated  a  "  spirit "  is 
given  off;  but  to  Priestley,  an  English  scientist,  is  given  the  honor 
of  making  the  discovery.  It  was  he  who  published  the  fact  to 
the  world,  and  his  observations  were  the  basis  of  a  correct  explana- 

21 


22 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


tion  of  combustion  or  burning.  As  long  as  a  discovery  is  unknown 
to  the  world  and  is  not  used,  it  is  of  little  value.  For  this  reason 
the  fact  that  Scheele,  a  Swedish  chemist,  recognized  oxygen  as  a 
definite  substance  before  its  discovery  by  Priestley,  is  of  historical 
interest  only  in  connection  with  an  account  of  Scheele's  life  and 
work. 

Priestley  was  a  clergyman  who  had  a  most  unusual  curiosity 
in  regard  to  nature;  he  took  up  natural  sciences  as  a  hobby,  and 
made  many  discoveries  of  the  greatest  importance  in  chemistry. 
His  theological  speculations,  which  were  recorded  in  many  volumes, 
have  been  forgotten,  but  the  results  of  his  study  of  natural  phe- 
nomena, which  he  carried  out 
in  the  spirit  of  the  amateur, 
are  among  the  foundation 
stones  of  the  science. 

Priestley  happened  to  have 
a  large  burning  lens,  and 
amused  himself  by  discover- 
ing the  effect  produced  when, 
by  means  of  it,  the  sun's  rays 
were  focused  on  various  ob- 
jects. In  one  of  his  experi- 
ments he  had  some  of  the  red 
oxide  of  mercury  floating  on 
the  surface  of  metallic  mer- 
cury contained  in  a  glass  tube 
(see  Fig.  3) .  When  the  heat 
was  concentrated  on  the 
oxide  of  mercury  it  slowly 
FIG.  3.  disappeared;  the  level  of  the 

mercury  fell  in  thetube,  which 

was  found  to  contain  an  air,  as  Priestley  said,  or  a  gas  as  we  say  to- 
day. This  gas  was  studied  by  him;  substances  burned  in  it  with 
great  brilliancy,  and  after  it  was  found  that  a  mouse  breathed  the 
new  gas  with  perfect  contentment,  Priestley  tried  it  on  himself. 
He  declared  he  was  refreshed  and  exhilarated,  but  deprecated 
the  use  of  the  substance  for  this  purpose,  and  said  he  believed  we 
should  be  content  with  the  air  which  nature  furnished  us. 

The  discovery  of  this  gas,  which  was  made  on  August  1,  1774, 


OXYGEN  23 

was  of  the  greatest  importance.  In  the  hands  of  Lavoisier,  a 
French  chemist,  it  put  chemistry  on  a  true  scientific  basis. 
Lavoisier  showed  that  this  gas  was  present  in  the  air,  and  that 
when  many  metals  rusted  or  were  heated  in  the  air,  they  united 
with  it.  He  showed,  further,  that  if  certain  of  the  products  so 
formed  were  heated,  they  decomposed  into  the  metal  and  the  gas. 
Lavoisier  burned  a  number  of  substances  in  the  gas,  and  found 
that  when  some  of  the  resulting  compounds  were  dissolved  in 
water,  acids  were  produced.  For  this  reason  he  called  the  gas 
oxygen,  deriving  the  name  from  two  Greek  words  signifying  acid- 
former. 

It  is  not  only  in  combustion  that  oxygen  plays  such  an  impor- 
tant part.  The  element  is  essential  to  life;  we  breathe  the  air  and 
the  oxygen  is  carried  by  means  of  the  lungs  and  blood  into  every 
part  of  our  body,  where  it  unites  with  the  materials  of  the  tissues. 
The  chemical  action  which  takes  place  produces  heat,  and  is  so 
regulated  that  the  temperature  of  the  body  remains  approximately 
constant  under  varying  outside  conditions.  In  the  rusting  of 
metals,  oxygen  takes  part;  it  brings  about  decay  in  dead  organic 
matter;  it  purifies  the  water  in  running  streams;  and  it  is  present 
as  a  constituent  of  all  living  things  and  in  many  mineral  sub- 
stances. Oxygen  forms  about  50  per  cent  of  the  atmosphere  and 
crust  of  the  earth;  eight-ninths  of  water  is  oxygen. 

It  is  evident  that  it  will  be  of  the  greatest  interest  to  study 
oxygen  in  considerable  detail.  We  shall  learn  how  to  obtain  it 
in  pure  condition,  study  its  properties  and  chemical  behavior,  and 
see  how  it  can  be  put  to  a  number  of  valuable  uses. 

22.  Separation  of  Oxygen  from  the  Air. — The  air,  we  shall  see 
later,  is  a  mixture  of  gases,  one-fifth  of  any  volume  of  it  being 
oxygen.  It  is  reasonable,  therefore,  to  assume  that  air  is  a  con- 
venient source  from  which  to  get  oxygen  directly.  For  a  long 
time  the  gas  could  not  be  obtained  in  this  way  since  there  was  no 
simple  method  known  to  separate  gases  from  one  another.  When 
a  method  was  devised  to  convert  air  into  a  liquid,  the  separation 
was  conveniently  effected.  A  liquid  boils  at  a  definite  tempera- 
ture. For  example,  water  is  converted  into  a  vapor  or  gas  at 
100°  when  the  temperature  is  measured  on  a  centigrade  ther- 
mometer (81);  wood  alcohol  boils  at  66°.  When  a  mixture  of 
the  alcohol  and  water  is  heated  to  boiling,  the  alcohol  changes 


24  INORGANIC  CHEMISTRY  FOR  COLLEGES 

into  a  gas  more  rapidly  than  the  water;  if  we  cool  the  vapor  which 
comes  off  first  to  the  temperature  at  which  it  liquefies,  we  obtain 
mostly  alcohol.  After  the  alcohol  has  boiled  away  the  liquid  left 
is  pure  water.  When  highly  compressed  air  is  cooled  to  a  low 
temperature  and  allowed  to  expand  rapidly,  it  changes  into  a 
liquid.  The  chief  constituents  of  air  are  oxygen  and  nitrogen. 
The  separation  of  the  two  substances  is  effected  by  boiling  as  in 
the  case  of  wood  alcohol  and  water.  Nitrogen  boils  at  a  lower 
temperature  than  oxygen,  consequently,  when  liquid  air  boils 
the  nitrogen  and  a  part  of  the  oxygen  pass  off  first.  After  the 
nitrogen  has  been  converted  into  a  gas,  what  is  left,  the  residue, 
is  oxygen.  The  method  has  become  a  practical  one  and  furnishes 
both  oxygen  and  nitrogen  for  commercial  use.  It  requires  elab- 
orate apparatus,  and  cannot  be  used  by  students  in  the  laboratory. 

23.  Preparation  of  Oxygen  by  Heating  Compounds  Contain- 
ing It.  (a)  Mercuric  Oxide. — Oxygen  is  conveniently  prepared 
by  heating  certain  substances  containing  it.  The  use  of  mercuric 
oxide  for  this  purpose  has  been  seen  to  be  of  historical  interest. 
The  preparation  can  be  conveniently  carried  out  as  follows: 
Some  of  the  oxide  is  placed  in  a  tube  of  hard  glass,  which  does  not 
soften  when  heated  in  a  flame  (Fig.  4).  The  tube  is  supported 
on  a  stand  and  connected  by  means  of  a  piece  of  rubber  tubing 
with  a  so-called  delivery  tube  bent  in  such  a  way  that  any  gas 
passing  through  it  can  be  collected  in  an  inverted  bottle  filled  with 
water  and  placed  over  the  end  of  the  tube.  As  the  gas  comes 
from  the  delivery  tube  it  forms  bubbles  in  the  water,  which  rise 
into  the  bottle  and  displace  the  liquid.  The  apparatus  by  means 
of  which  the  gas  is  collected  is  called  a  pneumatic  trough;  the 
value  of  its  use  in  collecting  gases  was  first  emphasized  by 
Priestley.1 

When  mercuric  oxide  is  heated  it  decomposes  into  the  ele- 
ments of  which  it  is  composed — oxygen  and  mercury.  The 
oxygen  being  a  gas  escapes  and  is  collected  in  the  bottle  over 
water.  As  the  heat  is  applied  the  mercury  also  passes  into  a  gas, 
but  changes  into  a  liquid  when  it  gets  out  of  the  region  of  the 

1  Some  of  Priestley's  original  apparatus  can  now  be  seen  in  the  National 
Museum  in  Washington.  He  was  driven  from  England  as  the  result  of 
religous  dissension,  and  settled  in  Northumberland,  Pennsylvania,  where 
he  founded  a  college. 


OXYGEN 


25 


flame;  it  settles  on  the  colder  parts  of  the  tube  in  drops,  which 
can  be  recognized  by  their  silver-like  color.  If  the  decomposi- 
tion is  continued  to  the  end,  the  red  oxide  disappears  entirely. 
The  gas  collected  can  be  shown  to  be  oxygen  by  inserting  into  it 
a  splinter  of  wood,  the  end  of  which  is  glowing.  When  a  piece  of 
wood  containing  sap  is  lighted,  it  burns  with  a  flame.  If  this  is 
extinguished  by  blowing  on  it  or  by  moving  it  rapidly  through 
the  air,  the  wood  continues  to  glow  for  some  time;  the  end  has 
the  appearance  of  a  hot  coal.  A  characteristic  property  of  oxygen 
is  that  it  causes  rapid  combustion,  and  when  a  glowing  splinter 


is  inserted  into  it,  a  brilliant  flame  is  produced  as  the  wood  burns 
rapidly.     This  effect  is  commonly  used  in  testing  for  oxygen. 

24.  (6)  From  Potassium  Chlorate. — Mercuric  oxide  is  an  expen- 
sive substance,  and  cannot  be  used,  therefore,  to  make  oxygen  on 
a  large  scale.  Potassium  chlorate  is  comparatively  cheap.  For  a 
long  time  it  was  the  principal  source  of  the  oxygen  which  was 
manufactured  for  commercial  use,  and  it  is  still  used  for  this 
purpose.  It  yields  oxygen  when  heated,  and  is  the  material  used 
in  the  laboratory  in  the  preparation  and  study  of  the  gas.  The 
apparatus  employed  is  like  that  just  described.  Potassium  chlorate 
is  familiarly  called  chlorate  of  potash;  it  is  used  as  a  wash  in  the 
case  of  sore- throat,  and  in  certain  tooth-pastes.  When  the  sub- 
stance is  heated  it  first  melts;  as  the  temperature  is  increased,  it 
begins  to  decompose,  and  bubbles  of  oxygen  rise  through  the 
melted  mass;  at  a  higher  temperature  the  decomposition  is  rapid. 


26  INORGANIC  CHEMISTRY  FOR  COLLEGES 

When  all  the  oxygen  has  been  set  free,  the  substance  left  is  a  white 
solid,  which  is  called  potassium  chloride.  Potassium  chlorate 
contains  three  elements  in  chemical  combination — potassium, 
chlorine,  and  oxygen.  The  chemical  change  which  takes  place 
when  it  is  heated  to  a  high  temperature  consists  solely  in  the  libera- 
tion of  the  oxygen  present.  The  resulting  compound,  potassium 
chloride,  contains  potassium  and  chlorine. 

25.  Catalytic   Action. — When   certain   substances   are   mixed 
with  potassium  chlorate,  the  rate  at  which  the  latter  decomposes 
is  increased.     This  can  be  shown  to  be  true  by  a  simple  experiment. 
A  test-tube  is  about  one-fourth  filled  with  potassium  chlorate, 
and  supported  by  means  of  an  iron  stand  and  clamp.     The  tube 
is  next  heated  to  the  temperature  at  which  decomposition  takes 
place  rapidly.     It  is  then  allowed  to  cool  until  the  evolution  of 
oxygen  appears  to  have  stopped  and  a  glowing  splinter  of  wood 
no  longer  bursts  into  a  flame  when  introduced  into  the  gas.     If, 
now,  a  small  quantity  of  powdered  manganese  dioxide — the 
amount  that  can  be  held  on  the  end  of  a  blade  of  a  pocket-knife — 
is  dropped  into  the  tube,  oxygen  is  rapidly  given  off.     This  can 
be  shown  by  again  testing  with  the  glowing  splinter;  it  will  burn 
brilliantly.     We  can  heat  manganese  dioxide,  the  substance  that 
produced  this  striking  result,  up  to  the  temperature  used  in  the 
experiment  and  it  yields  no  oxygen.     Further,   we  can  weigh 
carefully  some  manganese  dioxide,  use  it  to  cause  the  rapid  evolu- 
tion of  oxygen  from  potassium  chlorate,  recover  it  and  then  weigh 
it  again.     The  weight  remains  unchanged,  and  the  manganese 
dioxide  is  unaltered.     Many  cases  like  that  just  described  have 
been  studied.     A  variety  of  substances  are  known  which  bring 
about  reactions  that  do  not  appear  to  take  place  in  their  absence, 
although  the  substances  themselves  can  be  recovered  unchanged 
after  the  reaction  has  taken  place.     Such  substances  are  called 
catalytic  agents  or  catalyzers,  and  the  phenomenon  is  known  as 
catalysis  or  catalytic  action.     It  is  probable  that  the  part  played 
by  the  catalyzer  is  to  hasten  a  reaction  which  is  already  taking 
place  very  slowly.     If  a  substance  decreases  the  rate  of  a  reaction 
it  is  called  a  negative  catalyzer. 

26.  Many    natural    processes    are    brought   about   through 
catalytic  action.     The  energy  of  the  sunlight  is  stored  up  in  a 
growing  plant  as  a  result  of  chemical  actions  induced  by  chlorophyl, 


OXYGEN  27 

the  green  coloring  matter  in  leaves.  The  digestion  of  food  is 
accomplished  as  a  result  of  the  presence  of  catalytic  agents  in  the 
saliva,  the  gastric  juice,  and  other  fluids  of  the  body.  Catalysis 
has  been  studied  recently  in  great  detail,  and  the  process  has  been 
utilized  in  the  preparation  on  a  large  scale  of  many  substances  of 
industrial  importance.  We  shall  have  occasion  to  refer  to  a 
number  of  applications  later,  and  when  more  facts  are  at  command 
the  \vay  in  which  the  phenomenon  is  thought  to  take  place  will 
be  discussed. 

27.  Preparation  of  Oxygen  by  the  Action  of  Electricity  on 
Water. — We  have  seen  that  heat  is  an  important  agency  in  effect- 
ing chemical  change.     When  this  form  of  energy  fails,  the  desired 
result  may  often  be  produced  by  means  of  electricity.     Water  is 
a  chemical  compound  of  oxygen  and  hydrogen.     When  it  is  heated 
to  an  exceedingly  high  temperature  it  decomposes,  in  part,  into 
the  elements  of  which  it  is  composed;  but  as  these  are  both  gases 
they  cannot  be  separated  readily,  and  the  method  is  not  used  to 
prepare  oxygen.     When  electricity,  however,  is  passed  through 
water  it  is  broken  down  into  hydrogen  and  oxygen  without  the 
application  of  heat.     The  process  by  which  a  substance  is  decom- 
posed by  means  of  an  electric  current  is  called  electrolysis.     In 
order  to  render  the  water  a  conductor  of  electricity,  pure  water 
itself  offering  great  resistance  to  the  passage  of  the  current,  a 
small  amount  of  sulphuric  acid  is  added  to  it.     The  acid  acts  as 
a  catalytic  agent,  for  it  can  be  recovered  unchanged  after  a  large 
amount  of  water  has  been  decomposed.     The  experiment  can  be 
readily  carried  out  in  a  Hofmann  apparatus,  which  has  already 
been  described  (9).     When  the  decomposition  has  proceeded  for 
some  time  it  will  be  observed  that  the  volume  of  the  gas  which 
has  collected  in  one  tube  is  exactly  twice  that  in  the  other.     If 
the  gases  are  drawn  off  from  the  two  tubes  and  tested  separately, 
that  which  was  formed  in  the  smaller  amount  will  be  found  to  be 
oxygen.     It  will  cause  a  glowing  splinter  to  burst  into  flame. 
The  other  gas,  which  will  be  studied  in  detail  later,  is  hydrogen; 
it  will  extinguish  a  lighted  taper  thrust  into  it,  but  will  itself  burn 
with  an  almost  colorless  flame. 

28.  When  electricity  passes  through  anything  we  say  that  the 
place  from  which  the  current  flows  is  positive  and  the  place  to 
which  it  flows  is  negative;  this  is  done  as  a  matter  of  convenience 


28  INORGANIC  CHEMISTRY  FOR  COLLEGES 

in  describing  the  phenomena  produced  by  the  current.  In  the 
Hofmann  apparatus,  for  example,  the  current  flows  from  the 
electrode  in  one  tube,  through  the  water,  to  the  electrode  in  the 
other  tube.  The  current  passes  in  the  solution  from  the  pole  at 
which  oxygen  is  liberated  to  that  at  which  hydrogen  is  set  free. 
The  electrode  at  which  oxygen  is  set  free  is  called  the  anode; 
hydrogen  is  liberated  at  the  cathode.  If  we  disconnect  the  wires 
where  they  are  joined  to  the  apparatus  and  reverse  their  positions, 
the  poles  at  which  the  gases  are  evolved  will  be  reversed.  It  is 
evident  from  this  that  the  electricity  which  decomposed  the 
water  traveled  in  a  definite  direction;  it  is  what  is  called  a  direct 
current  of  electricity. 

Oxygen  and  hydrogen  are  produced  on  the  large  scale  for 
industrial  purposes  as  the  result  of  the  decomposition  of  water 
by  electricity. 

29.  Other  Methods  of  Preparing  Oxygen. — Oxygen  can  be 
prepared  from  a  number  of  compounds  containing  it.     A  sub- 
stance which  is  expensive,  but  is  sometimes  used  in  the  laboratory 
on  account  of  its  convenience,  is  sodium  peroxide.     There  are 
two  compounds  known  which  are  made  up  of  sodium  and  oxygen 
only.     They  differ  in  that  one  contains  a  greater  proportion  of 
oxygen  than  the  other;    the  one  containing  the  smaller  amount 
of  oxygen  is  called  sodium  oxide;    the  other  is  sodium  peroxide. 

The  latter,  which  is  a  white  powder,  is  formed  by  burning 
metallic  sodium  in  the  air.  When  sodium  peroxide  is  put  into 
water,  a  reaction  takes  place.  The  products  formed  are  sodium 
hydroxide,  which  dissolves  in  the  water,  and  oxygen,  which 
escapes.  The  presence  of  the  hydroxide  in  the  solution  can  be 
shown  by  the  fact  that  the  latter  will  change  the  color  of  red 
litmus  paper  to  blue.  Sodium  peroxide  which  has  been 
melted  and  poured  into  cans  while  in  the  liquid  condition  and 
allowed  to  solidify,  is  sold  under  the  commercial  name  "oxone." 
In  this  form  it  is  used  in  a  specially  constructed  generator  as  a 
convenient  source  of  small  quantities  of  oxygen. 

30.  A  number  of  substances  other  than  those  already  men- 
tioned yield  oxygen  when  heated  to  sufficiently  high  temperatures ; 
among  these  are  the  oxides  of  certain  metals.     The  temperatures 
to  which  these  compounds  must  be  heated  in  order  to  bring  about 
their  decomposition  into  oxygen  and  the  metal,  vary  widely. 


OXYGEN  29 

The  oxides  of  silver  and  gold  yield  oxygen  at  comparatively  low 
temperatures.  These  metals  do  not  react  readily  with  oxygen; 
they  do  not  rust;  they  are  said  to  be  chemically  inactive.  The 
oxide  of  magnesium,  on  the  other  hand,  cannot  be  decomposed  by 
heat  alone  at  the  highest  temperature  obtainable  on  the  earth. 
We  have  seen  that  magnesium  is  an  active  metal;  it  burns  rapidly 
in  the  air,  and  in  doing  so  gives  off  a  large  amount  of  energy  as 
heat  and  light.  Magnesium  oxide  is  a  very  stable  substance  on 
account  of  the  fact  that  the  elements  of  which  it  is  composed  lost 
so  much  of  their  chemical  energy  in  their  union.  Substances 
which  contain  a  large  amount  of  chemical  energy  that  can  be 
transformed,  are  active,  and  enter  readily  into  chemical  change. 
Those  which  are  formed  as  the  result  of  evolution  of  a  large 
amount  of  energy  are  more  stable. 

Certain  oxides  lose  a  part  of  their  oxygen  when  heated;  among 
these  are  copper  oxide  and  lead  dioxide.  Other  substances  which 
contain  a  large  proportion  of  oxygen  yield  the  element  at  a  high 
temperature;  but  a  consideration  of  these  is  deferred  until  they 
are  described  individually,  since  they  are  not  commonly  used  as 
a  source  of  the  gas. 

31.  Properties  of  Oxygen. — Oxygen  is  a  tasteless,  odorless, 
invisible  gas.  When  subjected  to  a  high  pressure  and  a  low  temp- 
erature it  changes  to  a  bluish  liquid,  and  at  a  still  lower  tempera- 
ture it  becomes  solid.  Oxygen  boils  at  — 182.5 °,1  and  melts  at 
—  227°;  the  liquid  or  solid  is  attracted  by  a  magnet.  One  liter 
of  oxygen  gas  when  weighed  under  standard  conditions  weighs 
1.429  grams.  Since  the  volume  of  a  gas  changes  when  the  pres- 
sure on  it  changes,  so-called  standard  conditions  have  been  defined 
in  order  to  facilitate  comparisons  between  different  gases;  the 
standard  temperature  is  0°  and  the  standard  pressure  is  that 
exerted  by  a  column  of  mercury  760  mm.  high.  This  pressure 
is  the  average  pressure  of  the  atmosphere  at  the  level  of  the  sea. 
Since  one  liter  of  air  under  these  conditions  weighs  1.293  grams, 
oxygen  is  slightly  heavier  than  air.  Oxygen  dissolves  sparingly 
in  water;  100  volumes  of  the  latter  dissolve  4  volumes  of  the  gas 
atO°. 

1  All  temperatures  given  refer  to  the  centigrade  scale,  which  is  described 
in  Section  81.  The  metric  system,  which  is  used  in  expressing  weights  and 
volumes,  is  explained  in  the  Appendix. 


30  INORGANIC  CHEMISTRY  FOR  COLLEGES 

32.  Chemical  Conduct  of  Oxygen. — The  behavior  of  oxygen 
when  it  enters  into  chemical  change  can  best  be  appreciated  by 
considering  some  striking  experiments.     A  number  of  large  jars, 
filled  with  the  gas,  are  provided.     Into  one  is  placed  a  glowing 
splinter;   at  once  the  wood  burns  rapidly  with  a  brilliant  flame. 
Into  another  is  introduced  a  bit  of  charcoal;  no  action  takes  place. 
The  charcoal  is  removed  and  heated  to  redness;  when  it  is  taken 
from  the  flame  and  allowed  to  remain  in  the  air  it  loses  its  bright- 
ness, and  quickly  cools.     It  is  heated  to  redness  again  and  thrust 
into  the  oxygen;    now  it  burns  with  a  brilliant  incandescence. 
The  experiment  is  repeated  with  sulphur.     When  cold  it  does  not 
react  with  the  oxygen.     When  ignited  it  burns  sluggishly  in  the 
air  with  a  blue  flame,  but  when  introduced  into  oxygen  a  flame 
giving  a  bright  blue  light  is  produced.     The  burning  of  phosphorus 
is  studied  next.     When  the  substance  has  been  warmed  it  burns 
in  oxygen  with  an  intense  white  light. 

These  experiments  show  that  the  substances  which  burn  in 
the  air,  burn  also  in  oxygen,  but  that  the  combustion  takes  place 
in  this  case  much  more  rapidly;  as  a  consequence  the  phenomena 
are  most  striking  and  brilliant. 

Some  substances  which  do  not  ordinarily  burn  in  the  air  do  so 
when  heated  and  placed  in  oxygen;  iron  furnishes  a  good  example. 
The  end  of  a  file  is  heated  to  redness  in  a  flame  and  put  into  a  jar 
of  oxygen;  it  burns  rapidly,  and  sufficient  heat  is  developed  to 
melt  the  metal;  and  bits  of  it  are  thrown  off  in  all  directions  like 
shooting  stars.  This  behavior  of  iron  is  utilized  in  one  form  of 
fireworks. 

33.  A  great  deal  can  be  learned  from  a  careful  study  of  these 
experiments.     We  observed  that  in  order  to  bring  about  the 
reaction,  in  each  case  the  substance  introduced  into  the  oxygen 
had  to  be  previously  heated.     The  temperature  to  which  a  sub- 
stance  must   be  heated   to   bring   about   its  rapid  union  with 
oxygen,  with  the  evolution  of  light,  is  called  its  kindling  tempera- 
ture.    The  word  kindling  as  here  used  has  the  same  significance 
as  it  has  commonly.     We  kindle  a  fire  by  heating  the  materials  to 
the  temperature  at  which  they  continue  to  burn  without  the 
assistance  of  heat  furnished  from  another  source.     The  tempera- 
ture at  which  substances  unite  rapidly  with  oxygen  is  markedly 
affected  by  their  physical  condition;  for  example,  a  metal  in  the 


OXYGEN  31 

state  of  a  very  fine  powder  has  a  kindling  temperature  much  lower 
than  the  same  metal  when  in  the  massive  condition.  Iron  can  be 
obtained  in  a  form  which  takes  fire  spontaneously  when  brought 
into  the  air. 

We  must  next  study  the  substances  formed  as  the  result  of 
burning,  or  combustion.  In  the  jars  in  which  the  wood  and  char- 
coal were  burned  we  find  nothing  that  we  can  see;  they  both 
contain  an  invisible  gas.  It  is  not  oxygen,  for  if  we  introduce 
into  it  a  glowing  splinter  it  is  at  once  extinguished.  A  new  sub- 
stance has  been  formed.  Both  the  wood  and  the  charcoal  con- 
tain the  element  carbon.  The  reaction  which  took  place  consisted 
in  the  chemical  combination  of  this  element  with  oxygen.  The 
product  is  an  oxide  of  carbon,  and  is  called  carbon  dioxide.  It  is 
formed  when  any  organic  substance  is  burned,  and  is  produced  in 
our  bodies,  and  is  present  in  exhaled  air. 

We  are  convinced  that  a  new  substance  has  been  formed  as 
the  result  of  burning  sulphur  in  oxygen  by  noting  the  odor  of  the 
gas  that  is  left.  If  we  breathe  it,  it  chokes  us  and  causes  violent 
coughing;  it  will  be  recognized  as  the  gas  formed  when  sulphur 
burns  in  the  air.  We  often  suffer  from  it  in  improperly  ventilated 
railroad  stations,  where  the  smoke  from  the  locomotive  is  allowed 
to  accumulate.  The  coal  contains  sulphur  and  when  it  is  burned 
the  gas  formed  from  the  latter  escapes  into  the  air.  The  sub- 
stance is  formed  as  the  result  of  the  union  of  sulphur  and  oxygen; 
it  is  called  sulphur  dioxide. 

The  jar  in  which  the  phosphorus  was  burned  contains  a  white 
powder;  it  is  phosphorus  pentoxide.  When  the  iron  burned  it 
was  converted  into  a  black  solid,  which  is  an  oxide  of  iron. 

In  all  cases  the  chemical  reaction  which  took  place  con- 
sisted in  the  direct  union  with  oxygen  of  the  substance  burned; 
and  an  oxide  was  formed  in  each  case.  All  compounds  which  con- 
tain oxygen  and  one  other  element  are  called  oxides.  The  prefix 
used  before  the  word  oxide  in  naming  compounds  refers  to  the 
amount  of  oxygen  present;  this  will  be  discussed  in  detail  later. 

A  careful  study  of  the  subject  has  shown  that  the  products 
which  are  formed  when  carbon,  sulphur,  phosphorus,  and  iron 
are  burned  in  oxygen,  are  the  same  as  those  produced  when  they 
are  burned  in  the  air.  This  fact  proves  to  us  that  air  contains 
oxygen. 


32 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


We  can  now  understand  as  a  result  of  our  experiments  what 
happens  when  a  substance  burns.  Burning,  or  combustion  as  it 
is  frequently  called,  consists  in  the  union  of  a  substance  with 
oxygen  with  the  simultaneous  production  of  light  and  heat.  The 
most  striking  chemical  property  of  oxygen  is  shown  in  the  act  of 
combustion;  we  commonly  state  this  by  saying  that  oxygen  sup- 
ports combustion.  The  substance  which  unites  with  oxygen  is 
said  to  be  oxidized;  the  process  is  called  oxidation. 

Oxygen  is  an  element.  No  one  has  ever  been  able  to  separate 
from  it  two  or  more  substances. 

34.  Slow  Oxidation. — Many  substances  which  burn  unite 
slowly  with  oxygen  without  the  production  of  light.  If  the  same 
product  is  formed  as  the  result  of  burning  and  of  slow  oxidation, 
the  same  amount  of  heat  is  given  off  in  the  two  cases.  When  the 
oxidation  is  slow  the  heat  is  evolved  at  such  a  rate  that  there  is 
no  appreciable  rise  in  temperature;  it  escapes  as  rapidly  as  it  is 

produced  and  the  substance  does  not  get 
hot  enough  to  emit  light.  An  instruc- 
tive experiment  will  help  to  make  this 
clear.  Some  steel  wool l  is  placed  in  a 
long  tube,  which  is  closed  at  one  end. 
The  tube  is  supported  in  an  upright 
position  by  a  clamp,  the  open  end  being 
placed  under  water  contained  in  a 
beaker.  (See  Fig.  5.)  The  tube  con- 
tains air  and  iron.  The  level  of  the 
water  in  the  tube  is  noted  carefully,  and 
the  apparatus  is  not  disturbed  for  some 
time.  At  the  end  of  a  day  or  two  it 
will  be  seen  that  the  iron  has  rusted, 
and  that  the  water  has  risen  in  the 
tube.  Iron  oxide  has  been  formed, 
and  as  the  oxygen  united  with  the 
iron  and  disappeared  as  a  gas,  the 
water  rose  to  take  its  place.  This  experiment  is  instructive,  as 

1  Steel  wool  consists  of  long  thin  shreds  of  the  metal  matted  together 
like  wool.  It  is  valuable  for  polishing  wood,  cleaning  floors,  etc.  It  is  used 
in  this  experiment  because  it  furnishes  a  large  surface  for  the  oxygen  to  act 
upon;  the  action  takes  place  more  rapidly  as  a  result. 


FIG.  5. 


OXYGEN  33 

it  proves  to  us  that  when  iron  oxide  is  formed  from  iron  and 
oxygen,  both  elements  disappear  as  such. 

35.  Spontaneous  Combustion. — Certain  substances  unite  so 
rapidly  with  oxygen  that  after  a  time  enough  heat  is  given  off  to 
raise  their  temperature  to  the  kindling  point,  and  light  is  produced 
as  they  burn.  For  example,  if  a  piece  of  phosphorus  is  left  in  the 
air,  it  soon  begins  to  give  off  a  white  smoke,  which  is  phosphorus 
trioxide,  produced  as  the  result  of  the  union  of  the  element  with 
oxygen.  The  chemical  action  takes  place  faster  and  faster;  heat 
is  produced  rapidly,  and,  finally,  the  kindling  point  is  reached 
and  the  phosphorus  bursts  into  flame.  Such  a  phenomenon  as 
this  is  called  spontaneous  combustion. 

Certain  substances  which  are  not  ordinarily  spontaneously 
combustible  exhibit  this  property  under  unusual  circumstances. 
Cotton  cloths  that  have  been  saturated  with  linseed  oil,  which  is 
used  in  making  paint,  and  left  undisturbed  for  some  time  in  a  pile, 
have  been  known  to  take  fire  spontaneously.  In  this  case  the  oil 
unites  with  oxygen  from  the  atmosphere;  as  cotton  is  a  poor  con- 
ductor of  heat,  and  the  cloth  inside  the  pile  is  protected  from  air 
currents  which  would  carry  away  the  heat  produced  as  the  result 
of  the  oxidation,  the  temperature  slowly  rises  until  a  flame  is 
produced. 

Cases  have  been  reported  of  the  spontaneous  combustion  of 
hay.  The  material  was  stored  in  a  damp  condition;  decomposi- 
tion took  place  as  the  result  of  oxidation;  and  as  hay  is  a  poor 
conductor  of  heat,  the  latter  accumulated  in  one  place  in  suffi- 
cient amount  to  finally  bring  about  active  combustion. 

36.  Incombustible  Substances. — Compounds  which  are  formed 
as  the  result  of  combustion  do  not,  in  general,  burn.     Substances 
like  stones,   bricks,   asbestos,   cement,   and  other  so-called  fire- 
proof materials,  all  contain  large  proportions  of  oxygen  and  cannot 
unite  with  more  of  the  element;  they  are  incombustible. 

37.  Uses  of  Oxygen. — Great  quantities  of  oxygen  are  con- 
sumed in  the  burning  of  coal  and  wood.     In  this  case  the  source 
is,  of  course,  the  air,  and  we  are  apt  to  overlook  the  fact  that  oxygen 
plays  such  an  important  part  in  our  happiness  and  comfort.     We 
have  no  particular  interest  in  the  carbon  dioxide  formed  when  the 
coal  burns,  and  let  it  escape  into  the  air.     What  we  desire  is  the 
energy  produced;  from  it  we  obtain  heat  and  power.     Coal  con- 


34  INORGANIC  CHEMISTRY  FOR  COLLEGES 

sists  essentially  of  carbon.  The  element  contains  chemical  energy. 
Oxygen  contains  chemical  energy.  When  the  two  unite  a  part 
of  the  total  energy  of  the  two  elements  is  transformed  into  heat. 
When  we  buy  coal  we  buy  chemical  energy. 

When  substances  burn  in  oxygen  a  higher  temperature  is 
attained  than  when  they  burn  in  the  air.  Only  about  one-fifth 
of  the  air  is  oxygen;  and  when  a  substance  burns  in  air  a  part  of 
the  heat  produced  is  used  to  warm  the  substances  present  which 
do  not  play  any  part  in  the  reaction;  as  a  consequence  the  temp- 
erature does  not  rise  so  high  as  when  pure  oxygen  is  used.  A 
number  of  uses  are  made  of  the  high  temperatures  produced  when 
substances  burn  in  oxygen.  Acetylene,  the  gas  which  was  formerly 
much  used  as  an  illuminant  on  automobiles,  produces  a  large 
amount  of  heat  at  a  high  temperature  when  it  burns  in  oxygen. 
The  form  of  apparatus  used  to  produce  a  flame  in  this  way  is 
called  an  acetylene  torch.  By  means  of  it  it  is  possible  to  cut  in 
a  short  time  masses  of  steel  that  yield  to  the  older  methods  only 
after  the  expenditure  of  much  time  and  labor.  The  flame  is 
slowly  drawn  across  the  metal;  at  the  high  temperature  produced, 
the  iron  burns  in  the  excess  of  oxygen,  the  oxide  formed  melts, 
flows  away,  and  soon  the  metal  is  severed.  In  this  way  it  is 
possible  to  cut  rails  and  girders  with  dispatch.  The  mass  of 
wreckage  which  resulted  from  the  collapse  of  a  great  steel  railroad 
bridge  was  cleared  away  with  the  aid  of  acetylene  torches.  By 
any  other  method  the  result  would  have  been  accomplished  only 
at  a  great  cost  in  time  and  labor.  This  is  another  interesting 
example  of  the  use  of  the  energy  produced  by  chemical  change. 
The  form  of  carbon  which  is  deposited  in  automobile  cylinders 
does  not  burn  in  air,  but  is  rapidly  converted  into  carbon  dioxide 
when  ignited  in  an  atmosphere  of  oxygen.  Use  is  made  of  this 
fact  in  cleaning  the  cylinders  of  gasoline  engines. 

Life  is  dependent  on  oxidation.  The  animal  quickly  dies  if 
deprived  of  oxygen.  In  certain  cases  of  extreme  weakness  in  a 
patient,  physicians  resort  to  the  use  of  oxygen.  If  breathing  is 
enfeebled,  it  is  advisable  to  inhale  pure  oxygen  rather  than  air; 
four-fifths  of  the  latter  is  an  inert  gas,  which  has  no  effect  on  the 
body. 

Oxygen  is  now  employed  with  great  success  to  restore  persons 
rendered  unconscious  by  noxious  gases.  Fire  departments  are 


OXYGEN  35 

supplied  with  an  apparatus,  called  a  pulmotor,  which  is  used  to 
introduce  oxygen  into  the  lungs  of  persons  who  have  been  suffo- 
cated in  a  fire. 

38.  Oxygen  in  Nature. — Having  studied  the  chemical  con- 
duct of  oxygen  we  can  now  understand  why  it  plays  such  an 
important  part  in  nature.  It  is  a  very  active  element;  it  forms 
compounds  with  nearly  all  of  the  known  elements.  When  the 
universe  was  being  formed  from  the  elements,  most  of  them  united 
with  oxygen;  and  we  find  oxides  of  many  elements  in  the  earth's 
crust.  Sand  is  made  up  chiefly  of  silicon  dioxide;  many  ores  of 
iron  and  other  metals  are  oxides;  and  water,  which  is  so  abun- 
dantly distributed  over  the  earth's  surface,  is  an  oxide  of  hydrogen. 

We  also  find  a  great  variety  of  minerals  which  contain  two  or 
three  elements  in  addition  to  oxygen;  limestone,  feldspar,  clay, 
and  many  other  substances  which  are  present  in  the  soil  belong  to 
this  class.  As  has  been  noted,  about  one-half  the  earth's  crust 
and  the  atmosphere  around  us  is  oxygen.  Not  only  the  earth, 
but  the  universe,  as  we  know  it,  contains  oxygen.  Vast  amounts 
of  it  are  present  in  the  sun — a  fact  which  is  discovered  by  the 
study  of  the  light  which  comes  to  the  earth  from  this  source. 
When  all  substances  are  heated  to  a  sufficiently  high  temperature 
they  give  off  light.  Each  element  under  these  conditions  emits 
light  which  is  made  up  of  certain  colors  the  combination  in  each 
case  being  characteristic  of  the  element.  With  the  aid  of  a  spe- 
cially devised  optical  instrument,  called  a  spectroscope  (615),  the 
colors  present  in  any  light  can  be  determined.  The  examination 
of  the  light  given  off  by  the  sun  shows  that  oxygen  is  present. 

In  the  decay  of  organic  matter  oxygen  plays  an  important 
part.  Through  the  agency  of  bacteria,  which  are  organisms  of, 
microscopic  size,  the  carbon  and  other  elements  present  are  con- 
verted into  oxides.  The  products  of  plant  and  of  animal  life  are 
transformed  in  this  manner,  and  find  their  way  eventually  into 
the  atmosphere  and  the  soil.  A  question  naturally  arises:  If 
such  great  quantities  of  oxygen  are  consumed  in  combustion  -and 
the  growth  and  decay  of  living  matter,  is  not  the  atmosphere 
losing  it  rapidly?  And  what  will  happen  when  this  element,  so 
necessary  to  life,  has  combined  with  other  things?  Nature  has 
provided  against  such  an  emergency.  When  plants  grow  they 
absorb  from  the  soil  and  the  air  the  waste  products  of  animal  and 


36  INORGANIC  CHEMISTRY  FOR  COLLEGES 

vegetable  life,  and  set  free  from  these  a  large  part  of  the  oxygen, 
which  in  this  way  passes  back  into  the  atmosphere.  In  this 
wonderful  transformation,  which  will  be  studied  more  fully  later, 
the  energy  of  the  sunlight  is  changed  into  chemical  energy  which 
is  stored  up  in  the  plant  and  the  free  oxygen. 

EXERCISES 

1.  State  the  effect  on  a  furnace  fire  of   (a)  opening  the  door  below  the 
grate  and  (6)  opening  the  door  above  the  fire.     Explain  why  the  observed 
results  occur. 

2.  What  occurs  when  a  wood  fire  is  fanned  or  blown  with  a  bellows,  and 
when  a  burning  match  is  placed  in  a  strong  draft  of  air  or  blown  upon  with 
the  breath?     Explain  the  reason  for  the  different  effects  produced  by  the 
same  cause. 

3.  When  water  is  decomposed  by  electricity  it  is  broken  down  into  hydro- 
gen and  oxygen.     How  do  you  think  these  gases  could  be  changed  into  water 
again? 

4.  How  could  you  tell  which  of  the  two  poles  of  an  electric  battery  to  be 
used  to  decompose  water,  is  the  positive  pole? 

5.  Sodium  oxide  and  sodium  peroxide  are  both  solids.     How  could  you 
tell  them  apart? 

6.  How  could  you  distinguish  potassium  chlorate  from  potassium  chloride? 

7.  Potassium  chlorate  is  used  in  fireworks.     Can  you  tell  why? 

8.  In  some  flash-powders  a  mixture  of  aluminium  and  potassium  chlorate 
is  used.     Explain  the  part  played  by  each. 

9.  Powdered  iron  is  used  in  a  form  of  fireworks  which  produces  a  mass  of 
brilliant  sparks.     What  substances  should  be  mixed  with  the  iron  to  make  it 
burn  rapidly? 

10.  Of  the  oxides  of  the  following  metals  which  ones  would  you  expect  to 
yield  oxygen  when  heated  in  a  gas  flame:  iron,  gold,  platinum,  silver,  copper? 

11.  Gold  and  platinum  are  used  to  fill  teeth.     Why  are  these  and  not  less 
expensive  metals  used? 

12.  If  you  were  asked  to  examine  jars  each  containing  one  of  the  following 
gases,  how  could  you  tell  what  substance  was  in  each  of  the  jars:   oxygen, 
air,  carbon  dioxide,  and  sulphur  dioxide? 

13.  If  you  were  asked  to  prepare  copper  oxide  and  zinc  oxide  how  would 
you  do  it? 

14.  If  you  were  given  a  sample  of  silver  oxide  how  could  you  prove  it  was 
not  an  element? 

15.  How  could  you  prevent  a  large  bulk  of  oily  waste  from  taking  fire 
spontaneously? 

16.  Why  do  we  turn  down  the  wick  of  a  lamp  that  is  smoking  to  stop  the 
formation  of  soot? 

J.7.  Why  does  not  gas  that  is  lighted  burn  back  into  the  pipe? 


OXYGEN  37 

18.  If  a  bit  of  wood  is  held  just  inside  the  top  of  a  lamp  chimney  it  does 
not  burn  with  a  flame.     Why? 

19.  Why  will  most  plants  not  grow  in  the  dark? 

20.  What  is  the  original  source  of  the  energy  produced  by  burning  wood? 

21.  When  iron  burns  in  oxygen  the  elements  lose  chemical  energy  which 
is  changed  into  heat.     When  the  iron  oxide  is  heated  with  carbon  (charcoal) 
the  metal  is  recovered  and  possesses  its  original  amount  of  energy.     Where 
did  the  latter  come  from  and  what  evidence  is  there  of  the  truth  of  your  con- 
clusion? 

22.  Water  contains  88.8  per  cent  oxygen.     How  many  grams  of  oxygen 
can  be  obtained  from  (a)  100  grams,  (6)  25  grams,  and  (c)  21.63  grams  of 
water?     How  many  liters  of  oxygen  are  obtained  in  each  case? 

23.  Potassium  chlorate  contains  39. 16  per  cent  of  oxygen,     (a)  How  much 
oxygen  is  contained  in  50.2  grams  of  potassium  chlorate?     (6)  How  much 
potassium  chlorate  must  be  taken  to  get  20  grams  of  oxygen? 


CHAPTER  V 
HYDROGEN 

Hydrogen,  which  we  have  seen  was  formed  along  with  oxygen 
as  the  result  of  the  electrolysis  of  water,  is  an  important  element. 
It  is  a  colorless  gas,  which  is  characterized  by  the  fact  that  it  is 
the  lightest  substance  known. 

39.  Occurrence  of  Hydrogen. — As  hydrogen  is  a  constituent 
of  all  living  things  it  is  widely  distributed  in  nature.     It  occurs  in 
the  free  condition — that  is,  not  in  combination  with  any  other 
element — in  the  gases  that  issue  from  volcanoes  and,  in  small 
quantities,  in  natural  gas.     Free  hydrogen  is  an  important  con- 
stituent of  the  gas  that  is  manufactured  from  coal  and  is  used  for 
the  production  of  light  and  heat.     Hydrogen  is  formed  as  the  result 
of  the  decomposition  of  organic  matter  brought  about  through 
the  action  of  certain  bacteria. 

The  occurrence  of  the  hydrogen  which  is  combined  with  other 
elements,  is  of  greater  importance  from  the  chemist's  point  of 
view.  It  constitutes  by  weight  one-ninth  of  water,  and  is  present 
in  all  substances  known  as  acids,  and  in  many  chemical  com- 
pounds. It  is  a  necessary  constituent  of  living  things. 

40.  Early  History  of  Hydrogen. — In  the  sixteenth  century 
Paracelsus  1  noted  the  fact  that  when  iron  was  treated  with  a  dilute 
acid  "  an  air  rises  which  bursts  forth  like  the  wind."     It  was  not 
until  1766,  however,  that  the  gas  was  really  discovered  and  studied. 
In  this  year  Cavendish,  an  English  physicist,  prepared  hydrogen 
by  the  action  of  hydrochloric  or  sulphuric  acid  on  zinc,  tin,  or 

1  Paracelsus  was  a  professor  of  medicine  in  the  University  of  Basle,  Switz- 
erland. He  was  the  first  to  discard  the  alchemical  doctrines  and  aims,  and 
to  put  aside  the  methods  based  on  magic  and  superstition,  which  were  used 
at  that  time  in  treating  disease.  He  studied  the  action  of  chemical  com- 
pounds on  the  body  and  used  them  as  drugs.  His  influence  was  great,  and 
men  turned  from  attempts  to  make  the  philosopher's  stone  to  a  more  prac- 
tical study  of  chemical  substances. 

38 


HYDROGEN  39 

iron.  He  isolated  the  gas  and  described  its  properties.  In  1781 
he  showed  that  when  hydrogen  burned  water  was  formed.  Some 
time  later,  Lavoisier  named  the  gas  hydrogen,  deriving  the  name 
from  the  Greek  words  signifying  water-former.  It  was  previously 
called  inflammable  air. 

41.  Preparation    of    Hydrogen,     (a)  By    the    Electrolysis    of 
Water. — We  have  already  seen  how  oxygen  and  hydrogen  can  be 
prepared  by  the  electrolysis  of  water  (27),  and  that  the  gases  are 
formed  in  the  ratio  of  one  volume  of  the  former  to  two  of  the  latter. 
We  also  discovered  that  hydrogen  burns  with  a  flame  which  is 
almost  invisible,  and  that  it  does  not  support  combustion.     If 
some  air  is  allowed  to  mix  with  the  gas  before  it  is  lighted,  the 
reaction  takes  place  with  the  production  of  a  characteristic  sound 
that  is  produced  as  the  result  of  an  explosion.     This  behavior  of 
the   gas — burning  when   pare   with   an   almost   colorless   flame, 
extinguishing  a  glowing  splinter,  and  exploding  when  mixed  with 
air  and  ignited — is  used  as  a  convenient  test  for  hydrogen. 

42.  (6)  Preparation  of  Hydrogen  by  the  Action  of  Sodium  on 
Water. — The  methods  which  have  been  described  up  to  this  point 
to  separate  oxygen  and  hydrogen  from  their  compounds,  are  based 
on  the  decomposition  of  the  latter  through  the  action  of  heat  or 
electricity;   energy  was  required  to  bring  about  the  change.     We 
are  now  about  to  see  how  chemical  energy  can  be  used  for  this 
purpose.     When  a  piece  of  sodium  is  put  upon  the  surface  of 
water,  a  violent  reaction  takes  place;  hydrogen  is  rapidly  evolved, 
and,  as  a  consequence  of  the  evolution  of  a  large  amount  of  heat, 
the  metal  melts,  and  finally  catches  on  fire  and  burns.     Sodium 
is  a  very  active  element,  and  contains  a  large  amount  of  chemical 
energy,   which  exhibits  itself  when  the  metal  is  brought  into 
contact  with  oxygen  or  certain  compounds  containing  oxygen. 
In  order  to  regulate  the  reaction  between  sodium  and  water  a 
simple  device  can  be  used.     A  small  piece  of  the  metal  is  wrapped 
in  copper  gauze;   when  this  is  put  into  water  the  weight  of  the 
gauze  causes  it  to  sink;    as  a  consequence,  the  sodium  cannot 
come  into  contact  with  the  oxygen  of  the  air,  and  cannot  burn. 
The  gas  evolved  is  collected  by  holding  a  tube  filled  with  water 
over  the  stream  of  bubbles  as  they  rise.     It  is  shown  to  be  hydro- 
gen by  inserting  into  it  a  burning  stick  of  wood;   the  gas  burns, 
and  the  flame  produced  by  the  burning  wood  is  extinguished. 


40  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  chemical  change  which  takes  place  in  this  striking  reac- 
tion is  this:  The  sodium  unites  with  the  oxygen  present  in  the 
water,  and  drives  out  a  part  of  the  hydrogen.  The  compound  of 
sodium,  hydrogen,  and  oxygen  which  is  formed  is  called  sodium 
hydroxide,  the  name  indicating  the  presence  of  the  three  elements. 
The  substance  is  a  white  solid  that  dissolves  in  water.  Many 
compounds  which  contain  hydrogen  and  oxygen  are  called  hydrox- 
ides; we  have,  for  example,  iron  hydroxide,  tin  hydroxide,  etc. 

The  metallic  hydroxides  which  dissolve  in  water,  such  as 
sodium  hydroxide,  show  the  characteristic  property  of  changing 
the  color  of  a  dye  called  litmus  from  pink  to  blue. 

43.  (c)  Preparation  of  Hydrogen  by  the  Action  of  Certain 
Metals  on  Water. — Elements  other  than  sodium  are  able  to  take 
the  oxygen  from  water  and  thus  set  free  hydrogen.  Whether  or 
not  any  given  element  can  do  this  is  determined  by  the  chemical 
energy  the  element  possesses.  Active  elements  like  magnesium, 
iron,  and  zinc  decompose  water.  In  these  cases,  however,  it  is 
necessary  to  apply  heat  to  bring  about  the  reaction,  or  to  hasten 
it  so  that  the  decomposition  of  the  water  takes  place  with  reason- 
able rapidity.  The  less  active  the  metal,  the  higher  it  must  be 
heated  before  reaction  occurs.  Magnesium  decomposes  water 
slowly  at  its  boiling-point,  zinc  at  a  higher  temperature,  and  iron 
must  be  heated  with  steam  to  dull  redness  before  hydrogen  is 
freely  evolved.  An  experiment  to  demonstrate  the  formation  of 
hydrogen  in  this  way  can  be  carried  out  easily.  A  supply  of  steam 
is  furnished  by  boiling  water  in  a  flask,  which  is  connected  with 
one  end  of  an  iron  tube  filled  with  iron  filings;  to  the  other  end  is 
joined  a  delivery  tube  which  leads  to  a  pneumatic  trough  (see 
page  42  for  a  drawing  of  the  latter).  The  iron  tube  is  supported 
in  a  furnace  supplied  with  a  number  of  gas  burners,  and  is  heated 
to  a  high  temperature.  As  the  water  boils  steam  passes  over  the 
iron  filings;  hydrogen  and  an  oxide  of  iron  are  formed,  and  the 
gas  escapes  through  the  delivery  tube  and  is  collected  over  water 
in  the  pneumatic  trough. 

The  most  inactive  metals  like  gold,  silver,  and  platinum  do 
not  decompose  water  at  any  temperature,  since  their  affinity  for 
oxygen  is  very  small  compared  with  that  of  hydrogen  for  oxygen. 
It  takes  a  large  amount  of  energy  to  separate  hydrogen  from 
oxygen  when  these  elements  exist  in  chemical  combination  as 


HYDROGEN  41 

water,   and   the  inactive  metals   are   not   able   to  furnish  this 
energy. 

44.  (d)  Preparation   of  Hydrogen    by   the   Action   of  Certain 
Metals  on  Acids. — We  shall  have  occasion  to  use  acids  repeatedly 
in  the  preparation  and  study  of  a  number  of  substances,  before 
these  acids  can  be  conveniently  discussed  in  detail.     For  a  correct 
understanding  of  the  chemical  reactions  into  which  they  enter,  it 
is  necessary  to  learn  something  in  a  general  way  about  a  few  of 
these  compounds.     A  characteristic  property  of  acids  with  which 
we  are  all  familiar  is  their  sour  taste.     Vinegar  contains  an  acid — 
acetic  acid — which,  when  separated  in  a  pure  condition,  is  a  color- 
less liquid  possessing  an  odor  which  we  recognize.     Sulphuric 
acid  is  a  heavy  oily  liquid,  and  is  sometimes  called  oil  of  vitriol. 
A  solution  of  the  acid  in  water  is  commonly  called  sulphuric  acid, 
and  the  substance  unmixed  with  water,  concentrated  sulphuric 
acid.     Sulphuric  acid  contains  the  elements  hydrogen,  oxygen, 
and  sulphur. 

Hydrochloric  acid  is  another  reagent  which  is  frequently  used. 
It  consists  of  a  solution  in  water  of  hydrogen  chloride,  which  is  a 
compound  of  the  elements  hydrogen  and  chlorine.  A  saturated 
solution  of  hydrogen  chloride  is  known  as  concentrated  hydro- 
chloric acid;  if  more  water  is  added  to  this  we  have  dilute  hydro- 
chloric acid. 

Hydrogen  is  present  in  all  acids;  and  all  acids  yield  hydrogen 
in  the  free  condition  when  they  are  treated  with  certain  metals. 
The  gas  is  most  conveniently  prepared  by  the  action  of  dilute 
hydrochloric  acid  on  zinc.  When  a  piece  of  the  metal  is  covered 
with  the  acid,  reaction  soon  begins,  bubbles  of  gas  appear  and  rise 
through  the  liquid,  and  the  solution  grows  warm.  As  the  tempera- 
ture rises  the  rate  at  which  hydrogen  is  set  free  increases  rapidly. 

45.  If  we  wish  to  collect  the  gas  for  study,  the  reaction  is 
carried  out  in  a  hydrogen  generator.     This  consists  of  a  bottle, 
closed  by  a  stopper,  through  which  two  tubes  pass.     One  of  these 
is  what  is  called  a  safety  tube,  or  thistle  tube  (Fig.  6) .     The  tube 
is  so  adjusted  that  one  end  of  it  nearly  touches  the  bottom  of  the 
bottle.     The  second  tube  passes  just  through  the  stopper  and  is 
bent  at  such  an  angle  that  it  can  be  conveniently  connected  with 
a  delivery  tube.     Some  zinc  is  placed  in  the  bottle,  the  stopper  is 
put  in  place,  and  acid  is  poured  in  through  the  safety  tube,  care 


42 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


being  taken  that  the  end  of  the  latter  is  beneath  the  surface  of  the 
liquid.  After  reaction  has  proceeded  for  some  time,  and  the  air 
which  was  in  the  bottle  has  been  driven  out,  the  delivery  tube  is 
placed  under  the  mouth  of  a  bottle  which  has  been  filled  with  water 
and  inverted  in  a  pneumatic  trough.  If  hydrogen  stops  coming 
off  as  a  result  of  the  exhaustion  of  the  acid,  more  of  the  latter  can 
be  added  through  the  safety  tube  without  admitting  air  to  the 
generator.  This  form  of  apparatus  is  simple  in  construction  and 
is  used  in  the  preparation  of  a  number  of  gases. 


FIG.  6. 

It  has  been  stated  that  hydrochloric  acid  is  a  compound  of 
hydrogen  and  chlorine.  When  zinc  reacts  with  the  acid,  the 
hydrogen  is  liberated  and  the  metal  unites  with  the  chlorine; 
the  compound  formed  as  a  result  is  called  zinc  chloride.  If  the 
water  which  is  present  is  boiled  away  the  compound  is  left  as  a 
white  solid. 

When  zinc  reacts  with  sulphuric  acid,  hydrogen  is  set  free, 
and  the  metal  unites  with  the  oxygen  and  ;ulphur  which  were 
present  in  the  acid;  the  compound  formed  in  this  case  is  called 
zinc  sulphate;  it,  too,  is  a  white  solid,  soluble  in  water.  Zinc 
and  acetic  acid  give  zinc  acetate  and  hydrogen. 

46.  A  number  of  metals  in  addition  to  zinc  liberate  hydrogen 
from  acids;  among  these  are  magnesium,  aluminium,  iron,  and 
tin,  all  of  which  are  more  or  less  active  metals.  Copper,  silver, 


HYDROGEN  43 

and  gold,  on  the  other  hand,  do  not  possess  this  power.  The  rate 
at  which  the  reaction  takes  place  in  the  case  of  several  metals 
varies  with  their  activity,  provided  other  conditions  are  the  same. 
This  can  be  readily  shown  by  putting  pieces  of  magnesium,  zinc, 
and  tin  of  equal  size  and  shape  into  hydrochloric  acid.  The  rate 
of  evolution  of  hydrogen  decreases  with  the  metals  in  the  order 
given. 

47.  The  rate  at  which  different  samples  of  any  one  metal  lib- 
erate hydrogen  from  an  acid  is  markedly  affected  by  the  purity  of 
the  material.     Zinc  of  the  highest  purity,  for  example,  reacts 
with  sulphuric  acid  with  extreme  slowness;    impure  zinc,  which 
contains  small  amounts  of  carbon  and  other  substances,  dissolves 
readily  in  the  acid.     The  impurities  act  evidently  as  catalytic 
agents.     A  simple  experiment  will  show  this  effect.     If  a  piece  of 
pure  zinc  is  put  into  dilute  sulphuric  acid  there  is  scarcely  any 
action.     If  now,  a  piece  of  platinum  is  placed  in  contact  with  the 
zinc,  hydrogen  is  evolved.     The  platinum  is  recovered  unchanged 
after  the  zinc  has  dissolved;  it  served  as  a  catalytic  agent.     Other 
metals  than  platinum  can  be  used  for  this  purpose.     The  follow- 
ing is  an  instructive  experiment:  Into  each  of  two  test-tubes  con- 
taining sulphuric  acid  is  placed  a  piece  of  zinc  which  reacts  slowly 
with  the  acid,  the  pieces  of  metal  used  being  approximately  of  the 
same  size  and  shape.     To  one  tube  is  added  a  few  drops  of  a  solu- 
tion of  copper  sulphate.     We  watch  carefully  what  happens:   A 
black  deposit  appears  on  the  zinc  in  the  tube  to  which  the  copper 
sulphate  was  added.     As  soon  as  this  is  formed  the  rate  at  which 
hydrogen  is  evolved  increases  markedly.     At  the  end  of  a  short 
time  the  reaction  which  has  been  catalyzed  takes  place  with  con- 
siderable rapidity,  whereas  the  other  proceeds  sluggishly,  if  at  all. 
The  black  deposit  formed  when  zinc  is  treated  with  a  solution  of 
copper  sulphate  is  metallic  copper;   the  color  of  the  latter  is  due 
to  the  fact  that  it  is  in  the  form  of  an  exceedingly  fine  powder.     In 
this  form  it  is  very  active  as  a  catalytic  agent  on  account  of  the 
fact  that  the  surface  of  the  metal  is  great,  and  there  are  very  many 
points  of  contact  between  the  copper  and  the  zinc.     A  small 
amount  of  copper  sulphate  is  often  added  to  hydrogen  generators 
to  increase  the  rate  at  which  the  gas  is  formed. 

48.  (e)  Preparation    of   Hydrogen   by   the    Action   of  Certain 
Metals  on  Bases. — We  have  learned  something  of  sodium  hydroxide 


44 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


in  the  study  of  the  action  of  sodium  on  water.  It  will  be  recalled 
that  hydrogen  and  this  compound  were  formed.  Sodium  hydrox- 
ide belongs  to  a  very  important  class  of  compounds  called  bases; 
Hydroxides,  as  the  name  implies,  contain  hydrogen  and  oxygen. 
From  some  of  these  hydrogen  can  be  readily  obtained  by  the 
action  of  the  more  active  metals.  We  can  trace  again  the  cause 
of  the  action  to  the  affinity  of  the  metals  for  oxygen;  they  seek 
out  this  element,  and  displace  the  hydrogen  united  to  it.  When 
aluminium  is  put  into  a  solution  of  sodium  hydroxide,  hydrogen  is 
set  free.  The  aluminium  unites  with  the  oxygen  and  sodium 
and  forms  a  compound  called  sodium  aluminate,  which  is  soluble 
in  water.  When  zinc  is  heated  with  solid  sodium  hydroxide  a 
similar  reaction  takes  place;  hydrogen  and  sodium  zincate  are 
formed. 

49.  Physical  Properties  of  Hydrogen. — Hydrogen  is  a  color- 
less, tasteless,  and  odorless  gas.  The  odor  observed  when  hydro- 
gen is  prepared  by  means  of  impure  metals,  such  as  commercial 
iron,  is  the  result  of  the  presence  of  impurities  in  the  gas.  Hydro- 
gen has  been  liquefied  and  solidified;  it  boils  at  —  252.5°  and 
melts  at  —  260°.  Hydrogen  is  the  lightest  substance  known;  one 
liter  of  it  under  standard  conditions  weighs  0.08987  gram.  It  is 

less    than    one-fourteenth 
as  heavy  as  air. 

If  we  wish  to  pour 
hydrogen  from  one  vessel 
into  another,  it  is  neces- 
sary to  hold  them  in  an 
inverted  position  (Fig.  7); 
as  the  bottom  of  one  vessel 
(a)  is  lowered  in  the  di- 
rection indicated  by  the 
arrow  6,  air  rushes  into  it 
and  the  hydrogen  forced 
out  ascends  into  the  second 
vessel  from  which  the  air 

is  expelled.  If  the  gas  in  c  is  now  tested  it  will  be  found  to  contain 
hydrogen.  The  lightness  of  the  gas  can  also  be  shown  by  filling 
soap  bubbles  with  it.  When  they  are  released  they  rise  rapidly 
in  the  air,  whereas  bubbles  filled  in  the  ordinary  way  with  air 


FIG.  7. 


HYDROGEN  45 

from  the  lungs  are  heavier  than  air  and  sink  to  the  ground.  If 
the  bubbles  containing  hydrogen  are  brought  into  contact  with  a 
flame  as  they  rise  into  the  air,  they  explode  with  a  loud 
noise. 

Hydrogen  is  adsorbed  by  many  metals,  the  amount  of  the  gas 
taken  up  in  any  case  being  affected  by  the  physical  condition  of 
the  metal.  Platinum  in  the  form  of  a  very  fine  powder  adsorbs 
about  fifty  times  its  volume  of  the  gas;  palladium  when  properly 
prepared  will  adsorb  over  500  volumes  of  hydrogen. 

50.  If  a  jar  containing  hydrogen  is  inverted  and  placed  mouth 
to  mouth  over  another  jar  containing  air,  and  the  two  are  allowed 
to  stand  undisturbed  for  a  few  minutes  and  are  then  examined,  it 
will  be  found  that  the  gases  have  mixed.     This  can  be  shown  by 
inserting  two  glass  plates  between  the  jars  and  then  separating 
them;  when  the  gases  are  tested  by  means  of  a  lighted  taper,  an 
explosion  occurs  in  each  case.     Since  hydrogen  is  so  much  lighter 
than  air  we  might  expect  it  to  remain  in  the  upper  vessel — to 
float  on  the  air  as  a  cork  does  in  water.     The  fact  that  it  does  not 
is  an  illustration  of  an  important  property  of  all  gases;  when  they 
are  brought  together  they  mingle  and  finally  a  uniform  mixture 
results;    this   property   is   called  diffusion.     Graham,   a   Scotch 
chemist,  who  studied  this  phenomenon  carefully,  discovered  a 
simple  law  which  expresses  the  relative  rates  at  which  gases  diffuse. 
The  rates  at  which  gases  diffuse  are  inversely  proportional  to  the 
square  roots  of  their  densities.     The  lighter  the  gas  the  faster  it 
diffuses.     One  liter  of  hydrogen  and  of  air  weigh  0.08987  gram  and 
1.293  grams  respectively.     The  rates  at  which  they  diffuse  are  as 
Vl.293  is  to  V0.08987  as  3.8  is  to  1.     Oxygen  is  sixteen  times  as 
heavy  as  hydrogen;  it  diffuses  one-quarter  as  rapidly  as  hydrogen. 

51.  The  fact  that  hydrogen  diffuses  more  rapidly  than  air  can 
be  shown  by  a  simple  experiment.     A  cup  of  unglazed  porcelain 
attached  by  means  of  a  stopper  to  a  glass  tube,  is  inverted  and  the 
lower  end  of  the  latter  placed  in  a  beaker  containing  water  (Fig. 
8).     A  jar  containing  hydrogen  is  next  placed  over  the  porce- 
lain cup.     The  hydrogen  diffuses  into  the  cup  through  the  pores 
of  the  latter  and  the  air  diffuses  out.    Air  is  forced  out  of  the  cup 
through  the  tube  and  rises  through  the  water  in  bubbles;   this 
takes  place  because  the  hydrogen  enters  the  cup  through  the 
pores  more  rapidly  than  the  air  passes  out  in  this  way.     The  gas 


46 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


accumulates,  as  a  consequence,  and  finally  forces  its  way  out 
through  the  water. 

52.  Chemical  Conduct  of  Hydrogen. — The  most  striking  chemi- 
cal property  of  hydrogen  has  been  exhibited  in  the  experiments 
which  have  been  described.  The  gas  burns  with  a  flame  that  can 
scarcely  be  seen.  If  any  solid  material  is  allowed  to  pass  through 

the  flame,  it  is  heated  to  such  a 
temperature  that  it  gives  off  light. 
A  convenient  way  to  do  this  is 
to  shake  together  near  the  flame 
two  blackboard  erasers  which 
contain  crayon  dust;  as  the  parti- 
cles of  the  latter  enter  the  flame 
they  become  incandescent  and 
give  off  yellow  light.  When  hy- 
drogen burns  it  unites  with,  oxy- 
gen and  forms  an  oxide.  It  will 
be  shown  later  that  this  oxide 
is  water. 

Hydrogen  does  not  unite 
readily  with  most  of  the  sub- 
stances which  burn  in  oxygen. 
When  a  burning  splinter  is  insert- 
ed into  the  gas  it  is  extinguished, 
because  wood  does  not  unite  with 
FIG.  8.  hydrogen.  The  gas  is  said  not 

to   support    combustion.     When 

this  experiment  is  carried  out,  the  hydrogen  itself  burns  at  the 
mouth  of  the  vessel  containing  it. 

When  hydrogen  and  oxygen  are  mixed  and  kept  at  the  ordi- 
nary temperature  no  apparent  reaction  takes  place  between  them. 
If  the  temperature  of  the  mixture  is  raised,  reaction  sets  in  very 
slowly  at  about  300°,  and  at  700°  the  gases  unite  with  explosive 
violence.  The  temperature  at  which  the  union  of  the  two  elements 
takes  place  is  affected  by  the  nature  of  the  substance  of  which  the 
vessel  containing  the  elements  is  made.  The  rate  of  the  reaction 
is  markedly  influenced  by  catalytic  agents.  An  experiment 
illustrates  this  clearly.  If  some  finely  divided  platinum  at  the 
temperature  of  the  room  is  introduced  into  a  stream  of  hydrogen 


HYDROGEN  47 

issuing  from  a  generator,  the  gas  immediately  bursts  into  a  flame. 
Platinum  for  this  purpose  is  prepared  by  soaking  some  asbestos,  a 
non-inflammable  mineral  substance,  in  a  solution  of  platinum 
chloride.  The  asbestos  is  then  heated  to  a  high  temperature; 
the  platinum  chloride  adhering  to  it  is  decomposed  by  the  heat, 
the  metal  is  deposited  as  a  very  fine  powder  or  sponge,  and  the 
chlorine  escapes  into  the  air.  Platinum  prepared  in  this  way  is 
an  active  catalytic  agent.  It  will  be  recalled  that  copper  in  a 
fine  state  of  division  was  used  to  catalyze  the  reaction  between 
zinc  and  sulphuric  acid.  As  the  catalytic  action  takes  place  at 
the  surface  of  the  metals,  it  is  evident  that  the  greater  the  surface 
the  greater  the  catalytic  effect. 

53.  When  a  mixture  of  hydrogen  and  oxygen  in  the  right 
proportions  is  ignited,  the  union  of  the  elements  takes  place  with 
explosive  violence.     This  fact  has  often  led  to  accidents  when 
attempts  have  been  made  to  light  the  gas  issuing  from  a  hydrogen 
generator.     When  the  acid  is  poured  upon  the  metal  and  action 
begins,  the  apparatus  is  full  of  air.     As  the  reaction  proceeds  the 
gas  forced  out  of  the  generator  is  a  mixture  of  hydrogen  and  air, 
the  proportion  of  hydrogen  increasing  as  the  metal  dissolves.     If 
a  flame  is  brought  in  contact  with  this  mixture  of  gases,  it  will  be 
ignited,  the  flame  will  travel  back  into  the  generator  and  a  violent 
explosion  will  occur;    and  the  flask  will  probably  be  shattered. 
An  experiment  to  illustrate  this  can  be  carried  out  with  safety  by 
covering  with  a  strong  box  the  hydrogen  generator  as  soon  as  the 
acid  is  introduced  into  it;  through  a  hole  in  the  side  of  the  box  is 
placed  the  tube  through  which  the  gases  issue.     A  lighted  burner 
is  placed  near  the  end  of  the  tube.     When  the  mixture  of  gases 
contains  its  constituents  in  the  right  proportions,  it  ignites;   the 
explosion  which  results  produces  a  loud  noise  and  shatters  the 
generator.     The  range  of  explosibility  of  mixtures  of  hydrogen 
and  air  is  wide;  if  such  a  mixture  contains  from  9.5  to  65  per  cent 
by  volume  of  hydrogen  it  will  explode  when  ignited. 

54.  An  explosion  is  produced  as  the  result  of  the  rapid  forma- 
tion of  gases  or  their  sudden  expansion.     When  a  mixture  of  oxy- 
gen and  hydrogen  is  ignited  a  large  amount  of  heat  is  generated 
in  an  incredibly  short  time.     Gases  increase  in  volume  when  they 
are  heated.     As  a  consequence,  the  water-vapor  formed  as  the 
result  of  the  union  of  the  oxygen  and  the  hydrogen  expands 


48  INORGANIC  CHEMISTRY  FOR  COLLEGES 

rapidly.  Since  the  heat  is  generated  so  quickly  its  full  effect  is 
produced,  and  there  is  not  time  for  it  to  be  lost  to  the  surround- 
ings. If  the  explosion  takes  place  in  an  open  vessel  with  a  wide 
mouth,  the  gases  can  expand  into  the  air,  and  the  vessel  is  not 
shattered.  The  air  is  set  in  rapid  motion  and  a  sound  is  produced 
as  a  result.  If  the  vessel  containing  the  exploding  mixture  is 
closed,  it  offers  a  resistance  to  the  tendency  of  the  gas  to  expand, 
and  pressure  is  produced.  Whether  or  not  the  containing  vessel 
is  destroyed  depends  upon  whether  it  is  strong  enough  to  resist 
the  pressure  generated  as  a  result  of  the  explosion.  The  vapor 
produced  from  substances  which  burn,  such  as  gasoline  and  ether, 
form  explosive  mixtures  with  air.  When  an  explosion  occurs  the 
hydrogen  in  such  compounds  unites  with  oxygen  to  form  water- 
vapor;  if  they  contain  carbon,  carbon  dioxide,  which  is  a  gas,  is 
formed. 

Hydrogen  unites  directly  with  a  number  of  other  elements, 
such  as  chlorine,  sulphur,  sodium,  etc.;  the  reactions  will  be  con- 
sidered when  these  elements  are  discussed. 

65.  Hydrogen  as  a  Reducing  Agent. — Hydrogen  has  a  strong 
tendency  to  unite  with  oxygen.  It  will  unite  not  only  with  free 
oxygen  as  we  have  seen,  but  also  with  oxygen  which  is  held  in 
chemical  combination  by  other  elements.  An  experiment  will 
illustrate  clearly  this  conduct  of  hydrogen.  A  hydrogen  generator 
is  attached  to  one  end  of  a  straight  tube  made  of  hard  glass,  which 
is  so  placed  that  it  can  be  heated  conveniently  by  a  number  of 
gas  burners;  to  the  other  end  is  joined  a  glass  tube  in  the  shape  of 
the  letter  U,  which  is  kept  cold  by  immersion  in  water.  Some 
copper  oxide  is  put  in  the  straight  tube,  and  hydrogen  from  the 
generator  is  passed  through  the  apparatus.  When  the  air  has 
been  expelled  the  copper  oxide  is  heated.  Reaction  soon  begins; 
the  black  copper  oxide  changes  to  metallic  copper,  which  can  be 
recognized  by  its  characteristic  color;  and  water  collects  in  the 
U-tube.  We  see  from  this  experiment  that  the  products  of  the 
reaction  between  hydrogen  and  copper  oxide  are  copper  and  water; 
it  consists  in  the  transfer  of  the  oxygen  from  one  element  to  the 
other.  The  copper  oxide  is  said  to  have  been  reduced.  Reduc- 
tion is  the  name  applied  to  the  process  by  which  oxygen  is  removed 
from  a  compound  in  which  it  is  present.  The  substance  which 
effects  the  change  is  called  a  reducing  agent.  One  of  the  most 


HYDROGEN  49 

important  chemical  characteristics  of  hydrogen  is  its  reducing 
power.  In  this  chemical  reaction  the  oxygen  that  is  removed 
from  the  copper  oxide  unites  with  the  hydrogen,  which  is  oxidized 
as  a  result.  In  general,  reduction  and  oxidation  take  place  simul- 
taneously; one  substance  is  reduced  and  the  reducing  agent  is 
oxidized. 

The  oxides  of  many  elements  can  be  reduced  by  hydrogen. 
The  ease  with  which  the  reduction  takes  place  is  determined  by  the 
activity  of  the  element  present  in  the  oxide.  In  the  case  of  the 
most  active  metals  the  transfer  of  oxygen  cannot  be  effected;  for 
example,  magnesium  oxide  and  aluminium  oxide  cannot  be 
reduced  by  hydrogen. 

Processes  of  oxidation  and  reduction  play  a  very  important 
part  in  chemistry;  many  metals  are  extracted  from  their  ores  by 
reduction.  Iron  is  obtained,  for  example,  by  heating  with  carbon 
oxides  of  iron  which  occur  in  nature;  the  carbon  serves  as  a 
reducing  agent  as  it  unites  with  the  oxygen  and  thus  sets  free  the 
iron. 

56.  Uses  of  Hydrogen. — The  uses  of  hydrogen  are  based  upon 
its  property  of  extreme  lightness,  upon  the  fact  that  it  contains  a 
large  amount  of  chemical  energy  which  is  transformed  into  heat 
when  the  gas  burns,  and  upon  its  reducing  power.  Most  of  the  hy- 
drogen used  commercially  is  obtained  as  a  by-product  in  the  manu- 
facture of  caustic  soda  and  chlorine  by  the  electrolysis  of  solutions 
of  sodium  chloride.  Hydrogen  is  used  in  balloons,  but  on  account 
of  the  cost  of  its  production  it  is  frequently  replaced  by  a  gas 
manufactured  by  a  special  process  from  coal. 

The  oxy-hydrogen  blow-pipe  is  a  form  of  apparatus  designed 
to  burn  hydrogen  in  oxygen  without  the  possibility  of  an  explosion. 
The  blow-pipe  consists,  in  brief,  of  a  tube  through  which  oxygen 
passes,  surrounded  by  a  second  tube  that  delivers  hydrogen.  The 
gases  mix  at  the  ends  of  the  two  tubes  as  they  pass  into  the  air. 
When  the  mixture  is  lighted  an  exceedingly  hot  flame  is  produced; 
its  temperature  is  said  to  be  2500°,  whereas  the  temperature  of  a 
Bunsen  flame  produced  by  burning  a  mixture  of  coal  gas  and  air  is 
about  1500°.  The  oxy-hydrogen  flame  is  used  when  a  source  of 
heat  at  a  high  temperature,  which  can  be  conveniently  manipu- 
lated, is  desired.  It  is  used  to  melt  and  work  platinum  and  other 
refractory  metals.  It  finds  an  important  application  in  the 


50  INORGANIC  CHEMISTRY  FOR  COLLEGES 

making  of  apparatus  from  quartz  or  silica,  a  substance  which 
resembles  sand  in  chemical  composition.  Quartz  cannot  be  melted 
in  an  ordinary  gas  flame.  Apparatus  made  from  it,  such  as  tubes, 
dishes,  retorts,  flasks,  etc.,  resists  the  highest  heat  commonly  used 
in  the  chemical  laboratory;  on  account  of  this  and  the  fact  that 
silica  apparatus  possesses  other  valuable  properties,  it  is  much 
used. 

57.  When  substances  which  do  not  burn  are  heated  to  a  high 
temperature  they  give  off  light,  the  brilliancy  of  which  increases 
very  rapidly  with  rise  in  temperature.     This  principle  is  made  use 
of  in  the  so-called  calcium-light  or  lime-light.     In  an  appropriate 
apparatus  a  piece  of  lime,  which  is  calcium  oxide,  is  heated  by  an 
oxy-hydrogen  flame.     Lime  serves  the  purpose  well,  because  at  the 
temperature  of  the  flame  it  does  not  melt;    it  emits  an  intense 
white  light,  which  was  formerly  much  used  in  projecting  stereopti- 
con  pictures  on  a  screen.     On  account  of  the  greater  ease  of  hand- 
ling electricity,  it  has  largely  replaced  the  calcium  light  for  this 
purpose,  although  the  steadiness  and  other  qualities  of  the  latter 
still  recommend  it.     Ordinarily  illuminating  gas,  which  contains  a 
large  amount  of  free  hydrogen,  is  used  in  place  of  the  pure 
gas,  on  account  of  the  ease  with  which  it  can  be  obtained.     The 
gases  used  in  connection  with  the  lime-light  are  stored  in  iron 
cylinders  under  pressure;   the  transportation  of  these  is  a  source 
of  inconvenience. 

Hydrogen  is  one  of  the  most  important  constituents  of  illumi- 
nating gas,  which  is  used  as  a  source  of  light,  heat,  and  power. 
The  hydrogen  is  obtained  in  one  form  of  this  gas  by  passing  steam 
over  hot  coal;  the  carbon  of  the  coal  removes  the  oxygen  from  the 
water,  and  sets  hydrogen  free.  The  reaction  in  this  case  is  similar 
to  that  of  the  action  of  steam  on  iron,  which  has  been  studied. 

58.  Hydrogen  is  much  used  as  a  reducing  agent  in  commercial 
chemistry.     Many  substances  which  are  transformed  into  dyes 
are  reduced  by  hydrogen  during  the  process.     Large  quantities 
of  hydrogen  are  used  in  making  solid  cooking  fats  from  liquid 
vegetable   oils,   such   as   cotton-seed   oil.     When  these   oils   are 
treated  under  pressure  with  hydrogen  in  the  presence  of  a  catalytic 
agent,  the  oil  and  the  hydrogen  unite  and  form  a  substance  which 
is  solid.     Finely  divided  nickel  is  the  catalytic  agent  used;   this 
metal  has  been  found  to  be  very  valuable  to  catalyze  reactions  in 


HYDROGEN  51 

which  hydrogen  is  involved.     The  direct  addition  of  hydrogen 
to  other  substances  is  called  hydrogenation. 

Large  quantities  of  hydrogen  are  used  in  the  manufacture  of 
ammonia.  In  this  process,  which  was  developed  in  Germany 
during  the  recent  war,  hydrogen  and  nitrogen  under  pressure  are 
brought  into  reaction  through  the  agency  of  a  catalyst.  From 
the  ammonia  so  prepared  nitric  acid  was  manufactured  for  use  in 
the  preparation  of  explosives. 

EXERCISES 

1.  Would  you  expect  (a)  water  to  be  decomposed  when  it  is  heated  with 
mercury  and  (6)  mercury  oxide  to  be  reduced  to  the  metal  when  heated  with 
hydrogen? 

2.  How  could  you  show  that  sour  milk  contained  an  acid? 

3.  Name  the  compounds  formed  when  iron,  zinc,  tin,  lead,  and  aluminium 
are  dissolved  in  (a)  hydrochloric  acid,  and  (6)  in  sulphuric  acid. 

4.  What  is  the  source  of  the  heat  produced  when  a  metal  dissolves  in  an 
acid? 

5.  How  could  you  distinguish  by  chemical  means  (a)  zinc  from  silver,  (6) 
magnesium  from  platinum,  (c)  tin  from  silver,  (d)  aluminium  from  zinc? 

6.  How  could  you  show  that  paper  contains  the  element  hydrogen? 

7.  What  is  the  source  of  the  hydrogen  that  is  present  in  combination  in 
a  piece  of  wood? 

8.  (a)  Why  would  you  expect  a  mixture  of  air  and  gasoline  vapor  to 
explode?     (6)  What  use  is  made  of  this  reaction?     (c)  Why  should  gasoline 
be  stored  in  a  place  where  there  is  a  free  access  of  air? 

9.  In  testing  samples  of  hydrogen  from  a  generator  to  find  out  if  air 
is  present,  the  tube  containing  the  gas  is  held  with  the  mouth  down  until  it 
has  been  ignited.     Why? 

10.  Give  examples  of  common  phenomena  which  are  produced  as  the 
result  of  the  fact  that  gases  diffuse. 

11.  If  a  porous  substance  like  silk  is  used  as  the  fabric  in  making  balloons 
to  be  filled  with  hydrogen,  the  gas  in  the  balloon  slowly  becomes  heavier  and 
its  lifting  power  is  reduced.     Why?     This  result  is  largely  avoided  by  oiling 
the  silk.     Why? 

12.  One  liter  of  air  weighs  1.29  grams  and  1  liter  of  hydrogen  0.09  gram. 
A  balloon  was  constructed  which  held  1000  cubic  meters  of  hydrogen,  and 
weighed  600  kilograms  when  inflated.     What  weight  would  the  balloon  just 
lift  from  the  eprth,  assuming  that  it  displaced  1000  cubic  meters  of  air? 


CHAPTER  VI 
THE  ATOMIC  THEORY.    CHEMICAL  EQUATIONS 

59.  Chemistry  advanced  very  rapidly  as  soon  as  the  balance 
was  used  as  an  aid  in  the  study  of  natural  phenomena.  Before 
this  time  fanciful  theories  had  been  put  forward  in  attempts  to 
explain  burning  and  rusting;  combustible  substances  and  metals 
that  rusted  were  supposed  to  contain  the  spirit  of  fire,  called 
phlogiston,  which  escaped  when  these  remarkable  changes  took 
place.  When  it  was  shown,  however,  that  a  metal  on  rusting 
increased  in  weight,  the  view  that  this  change  was  the  result  of 
the  escape  of  phlogiston  was  no  longer  tenable.  Lavoisier,  who 
studied  phenomena  of  this  kind,  showed  not  only  that  when  a 
metal  rusted  it  increased  in  weight,  but  that  this  increase  was 
equal  to  the  weight  of  the  oxygen  which  united  with  the  metal; 
the  discovery  of  oxygen  and  the  use  of  the  balance  thus  fur- 
nished the  true  explanation.  Lavoisier's  work  impressed  upon 
scientists  the  necessity  of  quantitative  measurements  in  the  study 
of  the  changes  that  take  place  in  matter;  it  resulted  further,  as 
we  shall  soon  see,  in  the  discovery  of  one  of  the  most  fundamental 
laws  of  nature — the  law  of  the  conservation  of  matter. 

We  have  learned  that  when  iron  burns  an  oxide  of  iron  is 
formed,  and  that  when  iron  rusts  an  oxide  of  the  metal  is  also 
formed.  A  close  examination  of  the  oxides  in  the  two  cases  brings 
out  the  fact  that  they  are  different  substances.  How  can  this  be? 
A  quantitative  study  of  the  two  reactions — that  is,  a  determina- 
tion of  the  quantities  of  the  iron  and  the  oxygen  that  react  in  the 
two  cases — will  answer  the  question.  The  two  compounds  differ 
as  the  result  of  the  fact  that  they  contain  iron  and  oxygen  in  differ- 
ent proportions — the  question  is  answered  by  an  appeal  to  the 
balance.  There  are  two  oxides  of  hydrogen;  quantitative  analy- 
sis shows  us  that  water  contains  a  smaller  percentage  of  oxygen 
than  the  second  oxide,  hydrogen  peroxide.  A  study  of  quanti- 

52 


THE  ATOMIC   THEORY.    CHEMICAL  EQUATIONS          53 

tative  relations  not  only  makes  clear  the  composition  of  substances, 
but  serves  many  practical  ends.  How  much  zinc  must  react  with 
hydrochloric  acid  to  produce  six  jars  of  hydrogen  to  be  used  for 
experimental  purposes?  How  much  cream  of  tartar  should  be 
mixed  with  a  certain  weight  of  cooking  soda  to  make  baking- 
powder?  Such  questions  as  these  can  be  answered  as  a  result  of 
the  study  of  the  quantities  of  substances  involved  in  chemical 
reactions.  Up  to  this  point  we  have  emphasized  the  qualitative 
side — the  side  which  has  to  do  with  the  different  kinds  of  matter 
that  undergo  change.  We  shall  now  see  that  our  knowledge  will 
be  greatly  increased  by  a  study  of  the  quantitative  aspect  of  these 
transformations. 

60.  Law  of  the  Conservation  of  Matter. — Lavoisier  and  others 
studied  many  chemical  changes  quantitatively;  as  a  result,  it  was 
found  that  the  sum  of  the  weights  of  the  substances  entering  into 
reaction  equaled  exactly  the  sum  of  the  weights  of  the  products  of 
the  change — no  matter  was  lost  or  gained.  This  conclusion  has 
been  repeatedly  confirmed.  Some  years  ago  the  German  chemist 
Landolt  carried  out  a  long  series  of  experiments  to  test  the  accu- 
racy of  this  conclusion;  he  worked  with  the  greatest  skill  and 
used  the  most  delicate  instruments  that  could  be  devised  for  de- 
tecting changes  in  weight.  Landolt  placed  the  substances  that 
were  to  interact  in  a  tube  shaped  like  an  inverted  letter  V;  one  sub- 
stance was  placed  in  one  arm  of  the  tube,  the  second  substance  in 
the  other.  The  tube  was  next  sealed  to  prevent  any  loss,  and 
was  carefully  weighed.  It  was  then  inverted;  the  substances 
contained  in  it  mixed,  and  the  chemical  reaction  took  place. 
After  some  time,  when  the  tube  returned  to  its  original  tempera- 
ture and  the  external  conditions  were  the  same  as  before  the 
reaction  took  place,  the  tube  was  weighed  again.  No  change  in 
weight  was  observed.  It  seemed  of  importance  to  test  with  the 
greatest  attainable  accuracy  the  conclusion  that  there  is  no  change 
in  the  total  weight  of  the  substances  which  are  undergoing  chemi- 
cal change,  for  this  conclusion  is  the  basis  for  all  quantitative 
study  in  chemistry.  This  fundamental  fact  is  summed  up  in  the 
law  of  the  conservation  of  matter,  which  is  stated  in  various  ways. 
One  statement  is  that  the  total  amount  of  matter  in  the  uni- 
verse does  not  change;  another  form  of  statement  is,  we  can 
neither  make  nor  destroy  matter — we  can  alter  its  form  only. 


54  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  first  expression  of  the  law  is  too  broad,  for  it  transcends  our 
experience;  a  better  way  to  express  it  is  to  state  that  the  total 
amount  of  matter  in  any  system  undergoing  change  remains  con- 
stant, and  by  system  we  mean  the  sum  total  of  that  particular 
matter  involved  in  the  change.  We  can  now  appreciate  defi- 
nitely what  is  meant  by  the  expression,  a  law  of  nature.  In 
science  a  law  is  a  general  statement  which  expresses  the  conclusions 
drawn  from  the  study  of  related  facts.  The  law  of  the  conserva- 
tion of  matter  is  a  statement  drawn  from  a  study  of  the  quantity 
of  matter  involved  when  matter  changes  its  form.  It  is  the 
result  of  human  experience  gained  through  observation  and  experi- 
ment. 

61.  Law  of  Definite  Proportions. — As  soon  as  the  importance 
of  the  quantitative  aspect  of  chemistry  was  recognized,  scientists 
undertook  a  study  of  the  quantitative  composition  of  chemical 
compounds.     After  exhaustive  investigations  and  extended  con- 
troversies with  other  chemists  Proust,  a  French  scientist,  estab- 
lished facts  which  warranted  the  conclusion  that  the  composition 
of  any  definite  substance  remains  constant.     This  conclusion  is 
expressed  in  the  law  of  definite  proportions,  which  states  that 
any  chemical  compound  contains  the  elements  of  which  it  is  com- 
posed in  a  definite  and  unvarying  proportion  by  weight.     As 
applied  to  a  particular  example  this  law  says  that  the  composition 
of  the  substance  water  is  fixed;    it  always  contains  the  element 
hydrogen  and  the  element  oxygen,  and  the  proportion  by  weight 
in  which  these  elements  are  present  never  varies;   all  samples  of 
pure  water  under  all  conditions  consist  of  one-ninth  hydrogen  and 
eight-ninths  oxygen. 

62.  Law  of  Multiple  Proportions. — It  has  been  stated  that 
there  are  two  oxides  of  hydrogen — water  and  hydrogen  peroxide — 
and  that  the  existence  of  these  two  compounds  is  the  result  of  the 
fact  that  hydrogen  and  oxygen  can  unite  in  different  proportions 
by  weight.     Similar  relations  exist  in  the  case  of  other  elements; 
there  are  two  important  oxides  of  carbon,  carbon  monoxide  and 
carbon  dioxide;  there  are  three  well-characterized  oxides  of  iron; 
and  the  element  nitrogen  unites  with  oxygen  in  five  different 
proportions.     The  English  chemist  John  Dalton,  whose  name  is 
one  of  the  most  famous  in  the  history  of  chemistry,  was  studying 
such  facts  as  these  when  he  made  a  very  important  discovery. 


THE  ATOMIC  THEORY.      CHEMICAL  EQUATIONS        55 

He  determined  the  different  proportions  by  weight  in  which  two 
elements  united  and  discovered  a  very  simple  relation  between 
them.  He  found  that  in  carbon  monoxide  the  carbon  and  oxygen 
were  united  in  the  relation  of  12  parts  by  weight  of  the  former  to 
16  parts  by  weight  of  the  latter;  in  carbon  dioxide  he  found  12 
parts  of  carbon  united  with  32  parts  of  oxygen.  It  wa  a  striking 
fact  that  12  parts  of  carbon  unite  with  16  parts  of  oxygen  to  form 
one  compound,  and  with  just  twice  as  much  oxygen  to  form  the 
other.  Dalton  analyzed  other  compounds  and  found  that  the 
figures  obtained  led  to  similar  results.  He  determined  the  differ- 
ent amounts  of  an  element  which  united  with  a  fixed  weight  of 
another  element  to  form  the  various  compounds  of  the  two  that 
could  be  made.  The  ratio  between  the  varying  amounts  was  not 
always  one  to  two  as  in  the  case  of  the  two  oxides  of  carbon,  but 
the  ratio  was  always  that  of  simple  whole  numbers.  Many  cases 
have  been  examined  since  Dalton  published  his  results.  Sulphur 
forms  two  oxides;  in  one  the  sulphur  and  oxygen  are  present  in 
the  ratio  of  32  parts  by  weight  of  sulphur  to  32  parts  by  weight 
of  oxygen;  in  the  other  the  proportion  is  32  of  sulphur  to  48  of 
oxygen.  The  weights  of  oxygen  which  unite  with  32  parts  of 
sulphur  are  32  and  48;  these  numbers  are  in  the  ratio  of  2  to  3. 
There  are  five  oxides  of  nitrogen.  If  we  determine  the  weights  of 
oxygen  which  unite  with  28  grams  of  nitrogen  in  each  case  to  form 
these  oxides,  we  shall  find  that  these  weights  are,  respectively,  16, 
32,  48,  64,  80— numbers  which  are  in  the  ratio  of  1,  2,  3,  4,  5. 

On  what  was  perhaps  insufficient  evidence  Dalton  enunciated 
what  he  called  the  law  of  multiple  proportions,  but  subsequent 
investigation  has  shown  that  the  law  summarizes  the  facts.  The 
law  of  multiple  proportions  states  that  when  two  elements  unite 
to  form  more  than  one  compound,  the  weights  of  one  of  these 
elements  which  unite  with  a  fixed  amount  of  the  second,  are  in 
the  ratio  of  small  whole  numbers. 

63.  The  Atomic  Theory. — Dalton  was  much  interested  in  the 
changes  that  take  place  in  the  weather;  he  made  daily  observa- 
tions and  recorded  them  for  many  years.  He  observed  that  the 
atmosphere  contained  varying  amounts  of  moisture  from  day  to 
day,  and  he  sought  to  explain  how  water-vapor  could  mix  with 
air.  Dalton  came  to  the  conclusion  that  this  would  be  possible  if 
air  were  composed  of  very  small  particles  which  moved  about 


56  INORGANIC  CHEMISTRY  FOR  COLLEGES 

freely  and  were  separated  by  spaces  of  considerable  size  as  com- 
pared with  the  size  of  the  particles.  He  considered  water-vapor 
as  made  up  in  the  same  way.  If  air  and  water-vapor  are  brought 
together  the  particles  can  mingle  and  form  a  uniform  mixture. 
He  called  these  small  invisible  particles  atoms.  When  the  law  of 
multiple  proportions  was  discovered,  Dalton  saw  at  once  that  his 
conception  of  atoms  offered  a  simple  explanation  of  the  remarkable 
facts  summarized  in  the  law.  The  law  of  definite  proportions 
could  also  be  explained.  Dalton  stated  that  all  substances  are 
composed  of  atoms,  and  that  when  these  unite  compounds  are 
formed.  Elements  are  made  up  of  atoms,  all  of  which  are  alike 
in  substance  and  weight.  Dalton  recognized  the  importance  of 
determining  the  characteristic  weights  of  the  atoms  of  the  several 
elements,  but  as  it  was  impossible  to  isolate  a  single  atom  and 
weigh  it,  he  took  upon  himself  the  task  of  determining  their  rela- 
tive weights.  We  shall  see  later  how  this  can  be  done.  Dalton 
proposed  the  atomic  theory  in  a  book  entitled  "  A  New  System  of 
Chemical  Philosophy,"  which  was  published  in  1808. 

64.  Let  us  take  the  present-day  conception  of  atoms  as 
developed  from  the  theory  of  Dalton  and  see  how  it  furnishes  an 
explanation  of  the  quantitative  relationships  observed  in  chemical 
action.  All  substances  are  made  up  of  exceedingly  small  particles 
called  atoms.  The  atoms  of  any  particular  element  are  alike  in 
substance  and  weight,  but  differ  in  these  two  respects  from  the 
atoms  of  all  other  elements.  Chemical  combination  consists  in 
the  union  of  two  or  more  atoms  to  form  what  is  called  a  molecule. 
The  smallest  particle  of  a  chemical  compound  is  a  molecule, 
since  chemical  compounds  contain  two  or  more  elements.  This 
comparatively  simple  conception  of  the  composition  of  matter 
has  been  the  guiding  principle  in  chemistry  for  over  a  hundred 
years.  While  at  first  it  was  a  mere  hypothesis — a  reasonable 
guess — which  was  put  forward  to  help  explain  a  few  important 
facts,  it  has  grown  in  importance  as  it  has  been  investigated  and 
has  been  applied  to  new  discoveries  made  as  the  science  grew. 
It  has  become  so  firmly  established  as  the  basis  of  all  science  that 
it  is  recognized  to-day  as  a  fact.  Experiments  of  great  brilliancy 
in  conception  and  ingenuity  in  execution  have  demonstrated  the 
presence  of  atoms,  which  have  been  weighed  and  counted.  When 
we  recognize  the  fact  that  the  amount  of  hydrogen  that  can  be 


THE  ATOMIC  THEORY.      CHEMICAL  EQUATIONS         57 

held  in  a  thimble  contains  about  the  number  of  atoms  expressed 
by  2  followed  by  20  ciphers,  we  can  begin  to  appreciate  to  what 
extent  man  can  go  in  unfolding  what  used  to  be  called  the  secrets 
of  nature. 

The  atomic  conception  of  matter  furnishes  a  reasonable  expla- 
nation of  the  facts  summarized  in  the  law  of  definite  proportions 
and  the  law  of  multiple  proportions. 

65.  Fact,  Law,  Hypothesis,  Theory,  Science. — In  the  growth 
of  knowledge  facts  are  first  discovered  through  observation  and 
experiment.  These  facts  are  next  sorted  out,  and  those  of  the 
same  nature  are  grouped  together  and  their  relation  one  to  the 
other  studied.  If  a  general  statement  can  be  made  which  sum- 
marizes a  large  number  of  facts,  a  law  is  discovered.  But  inquiry 
does  not  stop  here;  thinking  men  endeavor  to  find  some  reason 
for  these  laws;  they  picture  to  themselves  how  the  fundamental 
concepts,  matter  and  energy,  can  produce  the  observed  phenomena; 
and  the  result  is  a  hypothesis.  The  hypothesis  is  now  scruti- 
nized carefully;  is  it  in  accord  not  only  with  the  facts  and  laws 
which  led  to  its  production,  but  with  all  known  facts?  The  proc- 
ess known  as  deductive  reasoning  is  next  applied;  if  the  hypothe- 
sis is  true  what  consequences  follow?  Can  we  foretell  any  undis- 
covered facts?  The  conclusions  which  have  been  deduced  are 
next  tested;  if  they  prove  not  to  be  in  accord  with  the  facts  the 
hypothesis  is  rejected  or  modified.  When  a  hypothesis  has  been 
repeatedly  tested  in  this  way,  and  has  been  shown  to  be  in  accord 
with  all  known  facts  it  is  called  a  theory.  This  way  of  increasing 
knowledge  is  called  the  scientific  method.  The  sum-total  of  the 
facts,  laws,  and  theories  pertaining  to  any  particular  branch  of 
knowledge  is  called  a  science.  Science  has  been  defined  as  sys- 
tematized human  knowledge.  The  development  of  any  particular 
science  is  dependent  upon  the  extent  to  which  the  facts  considered 
in  that  science  have  been  systematized.  Some  sciences  are  largely 
descriptive,  that  is,  only  a  few  great  generalizations  have  been 
drawn  from  the  facts;  others  are  highly  developed.  The  more 
the  processes  of  mathematics  can  be  used  to  summarize  the  facts 
of  a  science,  the  more  developed  becomes  that  science.  Physics 
is  highly  developed;  we  shall  see  later,  for  example,  that  the 
behavior  of  gases  under  varying  conditions  of  volume,  pressure, 
and  temperature  can  be  expressed  by  a  very  simple  mathematical 


58  INORGANIC  CHEMISTRY  FOR  COLLEGES 

equation;  thousands  of  isolated  facts  are  indicated  by  the  expres- 
sion pv  =  RT,  when  the  significance  of  the  letters  used  in  the 
equation  is  understood.  Chemistry  was  for  a  long  time  a  descrip- 
tive science,  but  its  development  in  recent  years  has  been  rapid. 

66.  Atomic  Weights. — We  have  learned  that  when  two  ele- 
ments unite  to  form  a  compound  the  relation  between  the  weights 
of  the  two  is  definite.     These  relations  have  been  determined 
with  great  care.     It  has  been  found,  for  example,  that  16  parts 
by  weight  of  oxygen  unite  with  2  parts  by  weight  of  hydrogen, 
16  of  sulphur,  12  of  carbon,  65  of  zinc,  and  so  forth.     Applying 
the  theory  of  atoms  and  molecules  to  such  figures  as  these,  it  has 
been  possible  to  assign  numbers  to  the  various  atoms  which  repre- 
sent their  relative  weights;  the  number  assigned  to  each  element 
is  called  its  atomic  weight.     The  interesting  way  in  which  this  is 
done  will  be  described  later,  since  it  is  necessary  to  acquire  a  more 
extensive  knowledge  of  chemistry  before  it  can  be  thoroughly 
understood. 

Hydrogen  is  the  lightest  substance  known  and  its  atom  was 
taken,  therefore,  as  the  standard  of  weight;  the  atomic  weight 
of  hydrogen  is  1.  Since  an  atom  of  oxygen  weighs  sixteen  times 
as  much  as  an  atom  of  hydrogen,  its  atomic  weight  is  16.  Similarly 
when  we  say  the  atomic  weight  of  carbon  is  12  we  mean  that  an 
atom  of  carbon  is  12  times  as  heavy  as  an  atom  of  hydrogen. 

A  list  of  the  elements  with  their  atomic  weights  is  given  on  the 
inside  of  the  back  cover  of  this  book.  The  student  will  have 
occasion  to  use  these  numbers  repeatedly,  but  it  is  inadvisable  to 
attempt  to  learn  them  by  heart;  as  the  more  important  ones  are 
used  over  and  over  again,  their  values  will  be  remembered  without 
effort. 

67.  Symbols  and  Formulas. — The  use  of  letters  to  represent 
atoms  has  materially  simplified  the  method  of  recording  chemical 
reactions.     The  alchemists  used  astronomical  figures  to  represent 
the  elements;  the  symbol  O  or  >|<  for  the  sun  was  used  to  repre- 
sent gold,  the  crescent  moon  C  f°r  silver,  etc.     Dalton  employed 
geometrical  figures;    O  represented  hydrogen,  O  oxygen,  0  car- 
bon,  O  0  O  carbon  dioxide,  etc.     It  was  some  years  later,  in 
1811,  when  the  practical  suggestion  was  made  by  Berzelius  to 
indicate  the  elements  by  letters  taken  from  their  names.     0 
stands  for  oxygen,  H  for  hydrogen,  N  for  nitrogen,  etc.    When 


THE  ATOMIC  THEORY.    CHEMICAL  EQUATIONS  59 

the  initial  letter  of  two  or  more  elements  is  the  same,  a  second 
letter  taken  from  the  name  is  added  in  some  cases;  thus,  C  is  the 
symbol  for  carbon,  Ca  for  calcium,  and  Cd  for  cadmium.  In  the 
case  of  some  of  the  metals  the  symbol  is  taken  from  the  Latin 
name  of  the  element;  Sn  is  the  symbol  for  tin  and  is  derived 
from  the  word  stannum;  ferrum,  iron,  gives  us  the  symbol  Fe  for 
this  metal,  etc. 

68.  The  composition  of  a  compound  is  represented  by  a  chem- 
ical formula,  which  is  made  up  of  the  symbols  of  the  elements 
present  in  the  compound.     Used  for  this  purpose  the  symbol 
signifies  one  atom  of  the  element,  or  that  weight  of  the  element 
which  is  numerically  equal  to  its  atomic  weight.     We  can  use  any 
system  of  weights,  but  if  a  statement  is  not  made  to  the  contrary 
the  gram  is  the  unit  understood.     The  atomic  weight  of  carbon, 
for  which  the  symbol  C  is  used,  is  12;  the  atomic  weight  of  oxygen 
is  16  and  its  symbol  is  0.     When  these  symbols  appear  in  a  chemi- 
cal formula,  as  CO,  it  means  that  the  molecule  of  the  substances 
represented  by  the  formula  contains  one  carbon  atom  and  one 
oxygen  atom;    since  these  atoms  have  definite  weights,  twelve 
and  sixteen  respectively,  the  formula  indicates  that  in  this  sub- 
stance the  elements  are  united  in  the  proportion  of  12  parts  by 
weight  of  carbon  to  16  parts  by  weight  of  oxygen.     The  molecular 
weight  of  a  compound  is  the  sum  of  the  weights  of  the  atoms  it  con- 
tains.    If  grams  are  used  in  any  calculation  involving  the  above 
formula,  the  number  represents  what  is  called  a  gram-molecular- 
weight    or  mol,  that  is,  12  +  16  =  28  grams  of  the  compound. 
Likewise  28  pounds  is  a  pound-molecular-weight  of  it. 

69.  The  substance  to  which  is  assigned  the  formula  CO  is 
called  carbon  monoxide  to  distinguish  it  from  another  oxide  of 
carbon  the  molecule  of  which  is  made  up  of  one  carbon  atom  and 
two   oxygen   atoms — carbon   dioxide.     The   prefixes   indicate  in 
these  and  other  cases  the  number  of  atoms  of  any  particular  ele- 
ment present;  they  are  derived  from  Latin  words  signifying  one, 
two,  three,  etc. 

When  more  than  one  atom  of  an  element  is  present  in  a  com- 
pound, this  fact  is  indicated  in  its  formula  by  writing  a  number  to 
the  right  and  below  the  symbol  of  that  element;  the  formula  for 
carbon  dioxide,  which  contains  one  carbon  and  two  oxygen  atoms, 
is  expressed  thus:  C02.  A  number  placed  in  front  of  a  formula 


60  INORGANIC  CHEMISTRY  FOR  COLLEGES 

indicates  that  a  certain  number  of  these  molecules  is  considered; 
thus  2CO2  stands  for  2  molecules  of  carbon  dioxide. 

70.  Chemical  reactions  can  be  clearly  represented  by  means 
of  chemical  formulas.  For  reasons  which  will  be  given  later  we 
believe  that  oxygen  gas  is  made  up  of  molecules  each  of  which 
contains  2  atoms.  Atoms  are  represented  by  symbols;  mole- 
cules, which  contain  2  or  more  alike  or  unlike  atoms,  are  repre- 
sented by  formulas;  the  formula  for  oxygen  gas  is,  accordingly, 
62.  Likewise  the  formula  for  hydrogen  gas  is  H2.  Water  has 
been  shown  to  be  made  up  of  2  hydrogen  atoms  and  1  oxygen 
atom;  its  formula  is,  accordingly,  H^O.  With  the  aid  of  these 
formulas  we  can  express  the  fact  that  oxygen  unites  with  hydro- 
gen to  form  water;  it  is  represented  as  follows,  by  what  is  called  a 
chemical  equation: 

2H2  +  O2  =  2H20 

This  equation  signifies  that  2  molecules  of  hydrogen  unite  with 

1  molecule  of  oxygen  and  form  2  molecules  of  water.     Defi- 
nite weight  relations  are  also  indicated.     Since  hydrogen  is  the 
lightest  substance  known  the  weight  of  its  atom  may  be  taken  as 
the  unit  of  weight  of  atoms.     A  hydrogen  atom  weighs  1  unit  and 
an  oxygen  atom  weighs  16;  consequently,  a  molecule  of  hydrogen, 
H2,  weighs  2,  and  one  of  oxygen,  62,  weighs  32.     A  molecule  of 
water  weighs  2  +  16  =  18.     The  equation  states  that  2  mole- 
cules  of   hydrogen   weighing  2X2=4,  unite   with  1  molecule 
of  oxygen  weighing  32  and  form  2  molecules  of  water  weighing 

2  X  18  =  36.     To  express  this  clearly  we  can  rewrite  the  equation 
as  follows: 

2H2      +        O2  2H2O 

2(1  +  1)       (16  +  16)       2(1  +  1  +  16) 
4  +  32    =36 

These  numbers  are  relative  weights;  we  can  use  them  in  any  units. 

f  grams    j  f  grams    1 

Four  j  pounds  !•  of  hydrogen  react  with  32  \  pounds  >  of  oxygen  and 
I  ounces  J  I  ounces  J 

(grams    ^ 
pounds  \  of  water, 
ounces  J 


THE  ATOMIC  THEORY.     CHEMICAL  EQUATIONS          61 

When  symbols  are  put  together  in  the  way  shown  above  the 
result  is  called  an  equation,  because  the  number  of  atoms  on  one 
side  of  the  equality  sign  equals  the  number  on  the  other;  the 
equation  has  a  quantitative  significance.  We  could  express  the 
fact  that  hydrogen  and  oxygen  united  to  form  water,  in  the  fol- 
lowing way: 

_  H2  +  O2 


The  arrow  signifies  that  the  substances  represented  by  the  formulas 
to  the  left  of  it  change  to  the  substance  indicated  by  the  formula 
at  its  right.  This  qualitative  expression,  signifying  only  the  sub- 
stances involved,  lacks  the  definiteness  and  fullness  of  an  equation, 
which  furnishes  both  qualitative  and  quantitative  information. 
Some  authors  do  not  use  the  equality  sign  in  writing  chemical 
equations;  they  use  an  arrow  for  all  purposes  to  indicate  a  chemi- 
cal change.  In  this  book  the  equality  sign  and  arrow  will  have 
the  significance  which  has  just  been  explained. 

71.  The  Writing  of  Chemical  Equations.  —  In  order  to  write  a 
chemical  equation  we  first  set  down  the  formulas  of  the  substances 
which  interact,  separating  these  by  plus  signs  to  show  clearly 
each  formula.  We  next  write  in  the  same  way  the  formulas  of 
the  products  formed,  and  indicate  by  the  use  of  the  arrow  that  a 
transformation  has  taken  place.  The  reaction  between  zinc  and 
hydrochloric  acid  is  expressed  in  this  way  as  follows  : 

Zn  +  HC1  ->  ZnCl2  +  H2 

It  is  necessary  to  learn  the  formulas  of  substances  as  we  meet 
them.  Formulas  express  facts  which  have  been  established  as 
the  result  of  experiment.  Zinc  chloride  contains  the  metal  and 
chlorine  in  the  proportions  represented  by  the  formula;  and  it  is 
necessary  to  remember  that  the  formula  of  zinc  chloride  is  ZnCl2. 
We  shall  soon  see  that  if  we  know  this  formula  and  a  few  other 
facts  we  can  readily  write  the  formulas  of  a  large  number  of  com- 
pounds which  contain  zinc.  The  task  of  remembering  many 
formulas  is  not  so  great  as  might  appear  at  first;  there  is  a  beau- 
tiful system  underlying  chemical  combination,  which  once 
learned,  removes  the  necessity  of  relying  too  often  on  an  act  of 
mere  memory. 

Returning  to  the  chemical  equation  under  discussion,  we  see 


62  INORGANIC  CHEMISTRY  FOR  COLLEGES 

that  what  has  been  written  does  not  express  definitely  the  quanti- 
ties involved  in  the  reaction.  We  must  next  modify  the  expres- 
sion to  include  this;  we  balance  the  equation,  as  chemists  say. 
If  an  equality  sign  is  to  be  put  between  the  formulas  of  the  react- 
ing substances  and  those  formed,  the  resulting  equation  must  not 
violate  the  law  of  the  conservation  of  matter;  for  example,  we 
must  represent  as  much  hydrogen  on  one  side  of  the  sign  as  on 
the  other.  We  next  examine  each  symbol  and  see  if  the  same 
number  of  atoms  appear  on  the  two  sides  of  the  arrow.  There  is 
1  zinc  atom  on  the  left,  and  1  on  the  right;  no  change  is 
required  here.  There  is  1  hydrogen  atom  on  the  left  and  there 
are  2  on  the  right;  a  modification  is  necessary.  If  we  take  2 
molecules  of  hydrochloric  acid,  2HC1,  we  have  2  hydrogen  atoms; 
this  is  accordingly  done  and  a  2  is  placed  in  front  of  the  formula; 
the  expression  now  becomes, 

Zn  +  2HC1  ->  ZnCl2  +  H2 

We  continue  the  inspection  and  find  that  there  are  2  chlorine 
atoms  on  either  side  of  the  arrow.  The  number  of  atoms  of  each 
element  on  one  side  of  the  arrow  appears  to  be  the  same  as  that  of 
the  same  element  on  the  other  side.  We  run  through  the  inspec- 
tion of  each  element  a  second  time  to  see  that  the  placing  of  the  2 
in  front  of  HC1  has  not  altered  the  relations  which  were  examined 
before  this  change  was  made,  and  if  we  find  that  we  have  the  same 
number  of  zinc  atoms,  the  same  number  of  hydrogen  atoms,  and 
the  same  number  of  chlorine  atoms  on  the  two  sides  of  the  arrow, 
we  conclude  that  the  formulas  have  been  properly  balanced;  we 
change  the  arrow  to  an  equality  sign,  and  have,  then,  a  chemical 
equation : 

Zn  +  2HC1  =  ZnCl2  +  H2 

A  very  simple  example  has  been  taken  to  show  how  equations 
are  balanced;  the  consideration  of  a  more  difficult  one  will  be  in- 
structive. When  steam  is  passed  over  hot  iron,  hydrogen  and  an 
oxide  of  iron  are  formed.  Experiment  has  shown  that  the  oxide 
has  the  formula  FesCU;  this  is  a  fact  which  has  to  be  remembered. 
If  we  know  the  formula  of  water  and  that  of  hydrogen  we  can 
write  the  equation  for  the  reaction.  We  shall  first  set  down  the 
symbol  of  iron  and  the  formula  of  water,  since  these  two  substances 


THE  ATOMIC  THEORY.     CHEMICAL  EQUATIONS          63 

interact,  then  write  an  arrow,  and  next  the  formulas  of  the  sub- 
stances formed,  thus: 

Fe  +  H2O  ->  Fe3O4  +  H2 

Next  we  must  balance;  we  see  1  iron  atom  on  the  left,  and  3 
on  the  right;  place  a  3  before  the  symbol  of  iron;  the  expression 
is  now  as  follows: 

3Fe  +  H2O  -*  Fe3O4  +  H2 

If  the  iron  is  to  stay  balanced  we  cannot  change  the  number  of 
molecules  of  Fe304;  we  therefore  next  balance  the  oxygen,  the 
other  element  in  the  compound  containing  the  atoms  just  balanced. 
There  are  4  atoms  of  oxygen  on  the  right  of  the  arrow;  there 
must  be  4  on  the  left,  so  we  take  4  molecules  of  water. 
This  changes  our  expression  to  the  following: 

3Fe  +  4H20  -»  Fe3O4  +  H2 

The  number  of  molecules  of  water  must  not  be  changed  now,  for 
if  we  do  this,  the  balancing  of  the  iron  will  be  destroyed;  we 
therefore  balance  hydrogen  next.  There  are  8  atoms  of  this 
element  represented  to  the  left  of  the  arrow;  there  must  be  8 
on  the  other  side;  and  so  a  figure  4  is  put  in  front  of  the  formula 
for  hydrogen.  The  expression  now  becomes, 

3Fe  +  4H2O  ->  Fe3O4  +  4H2 

It  is  next  studied  to  see  if  all  the  atoms  are  balanced;  this  proves 
to  be  correct,  and  when  the  arrow  is  replaced  by  the  sign  of  equal- 
ity we  have  a  correct  chemical  equation.  The  advisability  of 
checking  up  the  work  by  a  re-examination  of  the  final  equation 
cannot  be  impressed  too  strongly,  for  such  an  examination  will 
detect  errors  that  might  have  been  made.  Another  point  is  to 
be  strongly  emphasized:  In  balancing  equations  the  numbers 
used  in  the  formulas  of  substances  to  express  their  composition, 
such  as  the  3  and  the  4  in  Fe304,  can  never  be  changed.  As  has 
been  pointed  out,  these  numbers  indicate  the  relative  proportions 
of  .the  elements  present  in  the  compounds.  For  example,  the  oxide 
of  iron  formed  by  the  action  of  steam  on  iron  always  has  the 
composition  indicated  by  the  formula  Fe304 — it  contains  iron  and 
oxygen  in  the  proportion  of  3  X  56  parts  by  weight  of  iron  to 


64  INORGANIC  CHEMISTRY  FOR  COLLEGES 

4  X  16  parts  by  weight  of  oxygen,  56  and  16  being  the  atomic 
weights  of  iron  and  oxygen  respectively.  Again,  if  in  balancing 
an  equation  containing  the  formula  H2O,  which  represents  the 
substance  water,  we  change  the  formula  to  H2(>4  to  bring  the 
oxygen  into  balance,  we  would  make  a  great  mistake,  for  there  is 
no  compound  with  the  formula  H2(>4,  and  if  there  were  it  would 
not  be  water.  To  state  the  fact  in  other  words,  in  balancing 
equations  all  we  can  do  is  to  put  what  numbers  we  please  before 
the  formulas  and  symbols.  This  change  means  that  we  use  and 
obtain  as  much  of  each  substance  as  is  required  by  the  law  of  the 
conservation  of  matter.  This  subject  has  been  gone  into  at  such 
length  on  account  of  its  importance,  and  the  fact  that  the  student 
is  apt  to  make  the  mistake  which  he  is  here  told  to  avoid. 

When  a  symbol  appears  in  the  formulas  of  two  substances  on 
the  same  side  of  an  equation,  balancing  is,  at  times,  a  little  more 
difficult.  We  will  write  the  equation  for  the  reaction  which 
takes  place  between  sodium  peroxide  and  water,  which  it  will  be 
recalled,  yields  sodium  hydroxide  and  oxygen.  In  order  to  be 
able  to  do  this  it  is  necessary  to  know  the  formulas  of  all  the  sub- 
stances involved;  these  are  given  below: 1 

Na2O2  +  H2O  -»  NaOH  +  02 

We  start  by  balancing  one  element,  usually  the  first  written  down, 
which  is  sodium  in  this  case.  Proceeding  in  the  way  illustrated 
above  we  get  the  expression, 

Na2O2  +  H2O  ->  2NaOH  +  O2 

Having  placed  a  2  in  front  of  the  NaOH  we  next  balance  the 
oxygen,  an  element  present  in  this  compound.  As  it  now  stands 
there  are  3  oxygen  atoms  to  the  left  of  the  arrow  and  4  to 
the  right;  the  oxygen  atoms  can  be  balanced  by  taking  2  mole- 
cules of  water: 

Na2O2  +  2H2O  ->  2NaOH  +  O2 

We  have  just  changed  the  number  of  water  molecules  to  balance 
oxygen,  so  we  now  attempt  to  balance  the  hydrogen,  the  second 
element  in  water.  There  are  4  atoms  to  the  left  of  the  arrow 
and  2  to  the  right.  To  balance  now  we  must  change  the  number 
1  The  symbol  for  sodium  (natrium)  is  Na. 


THE  ATOMIC  THEORY.     CHEMICAL  EQUATIONS          65 

in  front  of  the  formula  for  sodium  hydroxide  from  2  to  4,  and  this 
throws  out  of  balance  the  sodium  and  perhaps  the  oxygen.  It  is 
evident  that  the  equations  cannot  be  balanced  if  we  use  but  one 
molecule  of  Na2O2.  We  accordingly  place  a  4  in  front  of  the 
formula  NaOH,  and  balance  the  sodium  again  by  putting  2  before 
Na2O2.  Starting  with  2Na2O2  we  go  through  the  entire  process 
again  and  find  this  time  that  when  we  reach  the  last  element  it 
balances.  The  various  steps,  which  should  be  analyzed  carefully 
by  the  student,  are  as  follows: 

Na202  +  H2O  -»  NaOH    +  O2 

2Na2O2  +  H2O  -»  4NaOH  +  O2 

2Na2O2  +  2H2O  =  4NaOH  +  O2 

The  student  should  solve  the  problems  in  balancing  chemical 
equations  given  at  the  end  of  the  chapter.  The  method  is  used 
constantly  and  relieves  the  memory  of  much  effort;  it  is  not  neces- 
sary to  remember  the  number  of  molecules  entering  into  a  reaction. 
The  formulas  of  the  reacting  substances  and  those  formed  should 
be  kept  in  mind;  this  soon  becomes  an  easy  task,  since  these 
formulas  occur  again  and  again  in  other  connections. 

72.  Some  Simple  Chemical  Equations. — It  will  be  instructive 
to  express  in  equations  the  chemical  changes  which  have  been 
studied  thus  far. 

Magnesium  burns  in  oxygen  to  form  magnesium  oxide: 

2Mg  +  O2  =  2MgO 

When  mercury  1  and  iodine  are  rubbed  together  mercuric  iodide 
is  produced: 

Hg  +  I2  =  HgI2 

Iron  and  sulphur  when  heated  are  converted  into  ferrous  sul- 
phide: 

Fe  +  S  =  FeS 

1  The  symbol  for  mercury  is  Hg;  it  is  derived  from  the  Latin  name  of  the 
metal,  hydrargyrum,  which  means  liquid  silver.  The  symbol  for  iron,  Fe, 
comes  also  from  the  Latin  name  of  the  metal  which  is  ferrum;  K  is  the 
symbol  for  potassium,  which  was  called  kalium. 


66  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Oxygen  can  be  prepared  by  heating  mercuric  oxide: 

2HgO  =  2Hg  +  02 
by  heating  potassium  chlorate : 

2KC1O3  =  2KC1  +  3O2 
by  the  electrolysis  of  water : 

2H2O  =  2H2  +  O2 
and  by  the  action  of  water  on  sodium  peroxide: 

2Na2O2  +  2H2O  =  4NaOH  +  O2 

The  chemical  conduct  of  oxygen  was  shown  by  burning  in  it 
carbon,  sulphur,  phosphorus,  and  iron.  The  equations  for  these 
reactions  are  as  follows: 

C  +  02  =  C02 
S  +  O2  =  S02 
4P  +  5O2  =  2P2O5 
3Fe  +  2O2  =  Fe3O4 

Hydrogen  was  prepared  by  the  electrolysis  of  water: 

2H2O  =  2H2  +  O2 
It  was  formed  by  the  reaction  between  water  and  sodium: 

2Na  +  2H20  -  2NaOH  +  H2 
Iron  reacted  with  steam  to  produce  hydrogen: 

3Fe  +  4H2O  =  Fe3O4  +  4H2 

Hydrogen  was  also  prepared  by  the  action  of  certain  metals  with 
hydrochloric  acid,  HC1,  and  sulphuric  acid,  H2SO4,*  the  equa- 
tions for  these  reactions  are  as  follows: 

Zn  +  2HC1  =  ZnCl2  +  H2 

Fe  +  2HC1  =  FeCl2  +  H2 
Sn  l  +  2HC1  =  SnCl2  +  H2 
Zn  +  H2SO4  =  ZnSO4  +  H2 
Fe  +  H2SO4  =  FeSO4  +  H2 

1  Sn  is  the  symbol  for  tin;  it  is  derived  from  the  Latin  word  stannum. 


THE  ATOMIC  THEORY.     CHEMICAL  EQUATIONS          67 

Hydrogen  is  formed  when  zinc  and  aluminium  react  with 
sodium  hydroxide: 

Zn  +  2NaOH  =  (NaO)2Zn  +  H2 
2A1  +  6NaOH  =  2(NaO)3Al  +  3H2 

Hydrogen  burns  in  oxygen: 

2H2  +  O2  =  2H2O 
and  reduces  hot  copper  oxide: 

CuO  +  H2  =  Cu  +  H2O 

EXERCISES 

1.  Carbon  monoxide  contains  57.14  per  cent  oxygen,     (a)  How  many 
grams  of  oxygen  are  there  in  100  grams  of  carbon 'monoxide?     (6)  How  many 
grams  of  carbon  in  100  grams  of  carbon  monoxide?     (c)  How  many  grams  of 
oxygen  are  combined  with  1  gram  of  carbon  in  carbon  monoxide?     Carbon 
dioxide  contains  72.77  per  cent  of  oxygen,     (d)  How  many  grams  of  oxygen 
are  combined  with  1  gram  of  carbon  in  carbon  dioxide?     (e]  What  is  the 
relation  between  the  weight  of  oxygen  combined  with  1  gram  of  carbon  in 
carbon  monoxide  and  carbon  dioxide. 

2.  There  are  two  oxides  of  sulphur,  one  of  which  contains  50  per  cent  of 
oxygen  and  the  other  60  per  cent,     (a)  Calculate  the  weight  of  oxygen  united 
with  100  grams  of  sulphur  in  each  case.     (6)  What  relation  do  these  numbers 
bear  to  each  other?     (c)  Is  the  result  in  accord  with  the  law  of  multiple 
proportions? 

3.  Convert  the  following  into  balanced  chemical  equations: 

(a)  Al  +  HC1  -*  A1C13  +  H2. 
(6)  Mn  +  HC1  -*  MnCl2  +  H2. 

(c)  Sb  +  O2  -*  Sb2O3. 

(d)  Fe2O3  +  H2  -»  Fe  +  H2O. 

(e)  Ca  +  H2O  ->  Ca(OH)2  +  H2. 

(/)   Na2O2  +  H2SO4  -t  Na2SO4  +  H2O2. 
(g}  MnO2  -»  Mn3O4  +  O2. 
(h)  CuO  -»  Cu2O  +  O2. 


CHAPTER  VII 
CHEMICAL  CALCULATIONS 

73.  It  was  shown  in  the  last  chapter  that  the  conception  of 
atoms  and  molecules  and  a  study  of  the  quantitative  relations  in 
chemical  change  led  to  a  simple  method  of  expressing  in  the  form 
of  equations  the  reactions  which  take  place  between  substances. 
Since  these  equations  express  so  many  facts  in  a  very  brief  form, 
they  are  constantly  used  in  chemistry,  and  it  is  essential  that  their 
significance  and  the  uses  to  which  they  can  be  put  are  thoroughly 
understood. 

Let  us  take  a  simple  chemical  reaction  and  see  how  an  equation 
was  written  for  it  and  how  this  equation  can  be  used  when  once 
it  is  known.  When  hydrogen  is  passed  over  hot  copper  oxide  the 
products  formed  are  copper  and  water.  If  the  atomic  weights  of 
the  elements  are  known,  a  study  of  the  quantitative  relations 
involved  will  give  the  information  necessary  to  write  the  equation. 
The  atomic  weights  of  copper,  hydrogen,  and  oxygen  are  63,  1, 
and  16  respectively.  Experiments  showed  that  63  grams  of  copper 
unite  with  16  grams  of  oxygen  to  form  79  grams  of  copper  oxide, 
which,  consequently,  has  the  formula  CuO.  Experiments  also 
showed  that  2  grams  of  hydrogen  unite  with  16  grams  of  oxygen 
to  form  18  grams  of  water;  the  formula  of  the  latter  must  be, 
therefore,  IbO.  The  quantitative  relations  found  in  the  reaction 
between  copper  oxide  and  hydrogen  are  as  follows :  79  grams  of 
copper  oxide  react  with  2  grams  of  hydrogen  to  form  63  grams  of 
copper  and  18  grams  of  water.  These  facts  are  evidently  repre- 
sented by  the  following  equation: 

CuO  +  H2  =  Cu  -f  H2O 

It  is  important  for  the  student  to  see  clearly  how  the  weights 
given  lead  to  the  conclusion  expressed  in  the  equation.  The 

68 


CHEMICAL  CALCULATIONS  69 

results  of  the  several  experiments  have  been  recorded  in  this 
equation,  and  we  can  now  use  it  to  calculate  the  weight  relations 
involved  in  the  reaction.  It  is  important  to  note  that  all  chemi- 
cal equations  are  based  on  the  previous  study  of  the  weights  of 
the  substances  entering  into  the  reaction;  we  cannot  make  them 
up  out  of  our  heads,  or  guess  what  happens.  Chemical  formulas 
and  equations  are  a  kind  of  short-hand  with  which  we  can  record 
in  brief  form  many  facts.  All  the  facts  in  regard  to  the  reaction 
between  copper  oxide  and  hydrogen  stated  above  are  recorded  in 
the  equation  given ;  other  facts,  not  yet  brought  out,  are  indicated 
also.  When  one 'learns  how  to  read  this  chemical  short-hand, 
every  equation  is  a  source  of  much  information. 

74.  We  can  now  understand  why  we  are  able  to  use  chemical 
equations  to  calculate  the  weight  relations  involved  in  the  change 
expressed  by  the  equation.  Let  us  see  how  this  is  done.  How 
many  grams  of  hydrogen  will  be  required  to  react  with  100  grams 
of  copper  oxide?  We  will  rewrite  the  equation  for  the  reaction 
and  place  under  the  formulas  the  weight  relations  shown  by  the 
atoms  present.  We  must,  of  course,  know  the  atomic  weights, 
and  we  obtain  them  by  referring  to  the  table  on  the  inside  of  the 
back  cover  of  this  book.  Accurate  values  of  the  atomic  weights 
are  used  in  these  calculations.  In  order  to  make  the  discussion 
easier  to  follow,  round  numbers  have  been  used,  heretofore.  The 
accuracy  of  the  result  desired  determines  how  many  decimal 
places  should  be  retained  in  the  atomic  weights  and  the  calcula- 
tions. The  equation  for  the  reaction  described  above  is  as  fol- 
lows: 

CuO        +  H2  =     Cu    +  H20 

(63.57  +  16)  +  (1.008  +  1.008)       63.57  +  (1.008  +  1.008  +  16) 
79.57  2.016  63.57  18.016 

By  proceeding  in  the  way  indicated,  we  discover  the  relative 
weights  of  the  several  substances  which  take  part  in  the  reaction. 
From  these  numbers  we  can  calculate  by  simple  proportion  what 
weight  of  any  substance  is  involved  when  the  weight  of  any  other 
is  specified.  We  can  answer,  for  example,  the  question  stated 
above — how  many  grams  of  hydrogen  will  be  required  to  react 
with  100  grams  of  copper  oxide?  The  equation  tells  us  that 
79.57  grams  of  the  oxide  react  with  2  grams  of  hydrogen,  then 


70  INORGANIC  CHEMISTRY  FOR  COLLEGES 

CuO  :  H2         CuO  :  H2, 
79.57  :  2.016  =  100  :  x, 

79.57z  =  201.6 
*?ni  fi 

x  =        '     =  2.53  grams  of  hydrogen. 
/  y.o  i 

It  is  advisable  to  write  over  the  terms  of  the  proportion  the 
formulas  of  the  substances  involved;  neglect  to  do  this  often 
results  in  an  incorrect  statement  of  the  proportion.  Another 
form  of  stating  the  proportion  is  to  be  recommended;  it  reduces 
somewhat  the  amount  of  work  required.  The  .chemical  equation 
is  written,  as  before,  and  above  the  formulas  of  the  substances 
involved  in  the  particular  problem  are  placed  the  given  weight 
and  the  letter  x  to  indicate  the  weight  sought.  Beneath  these 
formulas  are  written  the  molecular  weights,  which  are  obtained 
by  adding  the  atomic  weights  of  the  atoms  in  the  separate  mole- 
cules. In  the  case  of  the  problem  just  solved  this  procedure 
would  lead  to  the  following  expression: 
100  gms.  x  gins. 

CuO      +    H2     =  Cu  +  H2O 

79.57          2.016 

From  these  numbers  we  make  a  simple  proportion  in  the  order  as 
written,  thus, 

100  :  x  =  79.57  :  2.016 

79.57s  =  201.6 

x  =  2.53  grams  of  hydrogen 

It  is  evidently  unnecessary  to  make  the  calculations  of  the  molec- 
ular weights  of  the  substances  not  involved  in  the  particular 
problem  being  solved. 

In  order  to  illustrate  the  method  further,  the  solutions  of 
other  problems  based  upon  this  reaction  will  be  indicated.  How 
many  grams  of  water  are  produced  when  25  grams  of  copper 
oxide  are  reduced  by  hydrogen? 

25  gms.  x  gms. 

CuO    +  H2  =  Cu  +    H20 

79.57  18.016 

25  :x  =  79.57:18.016 

79.57x  =  25  X  18.016 

x  =  5.66  grams  of  water 


CHEMICAL  CALCULATIONS  71 

How  many  grams  of  copper  can  be  obtained  by  reducing 
5  grams  of  copper  oxide  by  hydrogen? 

5  gms.  x  gms. 

CuO    +  H2  -          Cu         +  H2O 
79.57  63.57 

5  :  x  =  79.57  :  63.57 
79.57z  =  63.57  X  5 

x  =  3.99  grams  of  copper 

In  this  case  we  can  simplify  the  solution  somewhat.  As  all  the 
copper  comes  from  the  copper  oxide  we  can  set  down  the  problem 
as  follows: 

5  :  x 

CuO  :  Cu 
79.57    63.57 

The  rest  of  the  solution  is  as  given  above. 

As  a  further  example  of  the  method  one  more  problem  will  be 
solved.  How  many  grams  of  oxygen  can  be  obtained  by  decom- 
posing 10  grams  of  sodium  peroxide  with  water? 

10  gms.  x  gms. 

2Na202             +  2H2O  =  4NaOH  +  O2 

2(23  +  23  +  16  +  16)  16  +  16 

156  32 
10  :  x  =  156  :  32 
156z  =  10  X  32 

x  =  2.05  grams  of  oxygen. 

In  the  above  equation  2  molecules  of  sodium  peroxide  take  part 
in  the  reaction.  Each  molecule  contains  2  sodium  atoms  each 
weighing  23,  and  2  oxygen  atoms  each  weighing  16;  the  molecule 
weighs,  accordingly,  23  +  23  +  16  +  16  =  78.  Since  2  mole- 
cules are  involved,  this  number  must  be  multiplied  by  2. 

75.  The  same  method  of  calculation  can  be  used  to  determine 
the  percentage  of  an  element  in  a  compound  the  formula  of  which 
is  known.  The  words  per  cent  are  a  contraction  of  per  centum; 
10  per  cent  means  ten  per  hundred.  If  we  inquire  what  is  the 
percentage  of  oxygen  in  copper  oxide,  the  question  is  answered  by 


72  INORGANIC  CHEMISTRY  FOR  COLLEGES 

stating  how  many  grams  of  oxygen  are  contained  in  100  grams  of 
copper  oxide.     This  can  be  determined  as  follows: 


100  gms. 
CuO 
63.57  +  16 
79.57 
100:  x 
79.57z 

X 

x  gms. 
:         0 

16 
=  79.57  :  16 
=  1600 
=  20.09  per  cent. 

What  percentage  of  potassium  chlorate  is  oxygen? 

100  gms.  x  gms. 

KC1O3  :  3O 

39.10  +  35.46  +16  +  16+16       16  +  16  +  16 
122.56  48 

100  :  x  =  122.56  :  48 
122.56z  =  48  X  100 

x  =  39.16  per  cent 

76.  Calculations  Involving  Volumes  of  Gases. — It  is  often 
necessary  to  know  the  volume  of  a  gas  formed  in  a  given  reaction. 
This  can  be  done  with  the  aid  of  the  chemical  equation  for  the 
reaction  and  the  knowledge  of  the  weight  of  a  liter  of  the  gas. 
The  weight  of  the  gas  formed  is  first  determined  in  the  way  just 
described,  and  then  the  volume  of  this  weight  calculated.  A 
study  of  the  formulas  of  gases  from  the  standpoint  of  volumes 
leads  to  a  result  that  makes  the  method  of  calculation  simpler 
and  does  not  involve  a  knowledge  of  the  weight  of  any  particular 
gas.  Since  in  all  equations  the  molecular  formulas  of  the  gases 
are  used  and  these  represent  different  weights,  for  example,  Eb 
represents  2  grams  of  hydrogen  and  02  32  grams  of  oxygen,  it 
would  be  well  to  know  what  volumes  of  the  several  gases  these 
formulas  represent.  These  can  be  readily  calculated. 

What  is  the  volume  of  2  grams  of  hydrogen?  One  liter  of 
the  gas  has  been  found  by  experiment  to  weigh  0.09  gram  at  0°  and 
760  mm.  If  1  liter  weighs  0.09  gram  how  many  liters  weigh  2  grams? 

1  :  0.09  =  x  :  2,     x  =  22 A  liters 


CHEMICAL  CALCULATIONS  73 

The  formula  for  oxygen  is  62  and  its  molecular  weight  is 
2  x  16  =  32.  What  is  the  volume  of  32  grams  of  oxygen?  One 
liter  of  oxygen  weighs  1.429  grams  at  0°  and  760  mm.: 

1  :  1.429  =  x  :  32,  x  =  22 .4  liters 

If  we  calculate  in  this  way  the  volume  of  a  gram-molecular-weight 
(68)  of  any  gas  the  result  is  always  22.4  liters.  We  shall  learn 
later  (Chapter  XXIV)  the  reason  for  this  striking  fact.  We  see, 
then,  that  the  formulas  of  gases  have  a  definite  significance  in 
regard  to  volume  as  well  as  to  weight.  The  formula  CO2  assigned 
to  carbon  dioxide  means,  as  we  have  learned,  that  the  molecule 
of  the  gas  contains  1  atom  of  carbon  and  2  atoms  of  oxygen, 
and  since  the  atomic  weight  of  carbon  is  12  and  that  of  oxygen 
16,  it  means,  further,  that  12  grams  of  carbon  are  united  with 
2  X  16  =32  grams  of  oxygen.  We  see  now  that  44  grams  of 
CO2  (12  +  2  X  16=  44),  which  is  1  gram-molecular-weight,  has  a 
volume  of  22.4  liters. 

We  are  now  in  a  position  to  calculate  the  volumes  of  gases 
found  in  any  chemical  reaction.  The  method  is  illustrated  by 
the  following  example.  What  volume  of  hydrogen  is  formed 
when  10  grams  of  zinc  dissolve  in  hydrochloric  acid?  We  write 
the  equation  as  before, 

10  gms.  x  liters 

Zn      +  2HC1  =  ZnCl2  +  H2 
65.37  22.4  liters, 

but  place  below  the  formula  for  hydrogen  the  volume  of  the  gas 
formed  and  use  this  in  the  proportion  instead  of  the  weight  in 
solving  the  problem, 

10  :  x  =  65.37  :  22.4  liters 
x  =  3.42  liters. 

In  the  following  problem  the  question  is  put  in  a  different 
way  but  the  method  is  the  same.  What  weight  of  zinc  is  required 
to  liberate  5  liters  of  hydrogen? 

x  gms.  5  liters 

Zn      +  2HC1  =  ZnCl2  +  H2 
65.37  22.4  liters 

x  :  5  =  65.37  :  22.4 
x  =  14.6  grams. 


74  INORGANIC  CHEMISTRY  FOR  COLLEGES 

In  solving  problems  in  this  way  the  volume  is  always  22.4  liters 
for  each  gram-molecular-weight  whatever  the  gas  may  be.  If 
3  molecules,  for  example,  appear  in  the  equation,  then  the  volume 
of  the  gas  is  3  times  22.4  liters.  In  the  case  of  the  decomposition 
of  potassium  chlorate  the  equation  is  as  follows: 

2KC1O3  =  2KC1  +         3O2 

2[39  +  35  +  3(16)]  3(22.4  liters) 

244  67.2  liters 

77.  The  volume  of  1  gram-molecular-weight  of  a  gas,  namely, 
22.4  liters,  is  called  the  gram-molecular-volume.  Since  this  is  a 
constant  quantity,  it  is  possible  to  see  at  a  glance  the  relation  by 
volume  in  which  gases  interact,  if  the  equation  for  the  reaction  is 
known.  In  the  case  of  the  union  of  hydrogen  with  oxygen  to 
form  water: 

2H2         +          O2       =        2H2O 
2(22.4  liters)      22.4  liters       2(22.4  liters) 

if  the  temperature  is  such  that  the  water  formed  is  a  gas — i.e., 
steam — we  see  that  2  times  22.4  liters  of  hydrogen  unite  with 
22.4  liters  of  oxygen  to  form  2  times  22.4  liters  of  steam,  which 
means  that  2  volumes  of  hydrogen  react  with  1  volume  of  oxygen 
to  form  2  volumes  of  steam.  The  relation  between  the  volumes 
of  the  gases  involved  in  the  reaction  is  the  same  as  that  between 
the  number  of  molecules.  If  we  take  as  a  second  example  the 
equation, 

N2  +  3H2  =  2NH3 

which  expresses  the  union  of  nitrogen  with  hydrogen  to  form 
ammonia,  which  is  a  gas,  we  see  that  1  molecule  of  nitrogen  unites 
with  3  molecules  of  hydrogen  to  form  2  molecules  of  ammonia 
and  that  1  volume  of  nitrogen  unites  with  3  volumes  of  hydrogen 
to  form  2  volumes  of  ammonia.  The  relation  holds,  of  course, 
whatever  unit  of  volume  is  used — liters,  pints,  gallons,  bar- 
rels, etc. 

Since  1  gram-molecular-weight  of  all  gases  occupies  22.4  liters 
we  can  calculate  from  the  formula  of  a  gas  what  1  liter  of  it  weighs. 
For  example,  the  formula  of  sulphur  dioxide  is  S02.  Since  the 
atomic  weight  of  sulphur  is  32  and  that  of  oxygen  16,  the  molecu- 
ular  weight  of  sulphur  dioxide  is  32  +  (2  X  16)  =  64,  Sixty- 


CHEMICAL  CALCULATIONS  75 

four  grams  of  the  gas  have  the  volume  22.4  liters,  therefore,   1 
liter  weighs  64  grams  -5-  22.4  =  2.85  grams. 

78.  By  making  use  of  the  fact  that  the  volumes  of  1  gram- 
molecular-weight  of  all  gases  are  the  same,  we  can  readily  tell  by 
inspection  of  the  formulas  of  two  gases  which  is  the  heavier.  Which 
is  heavier,  nitrogen  or  oxygen?     The  molecular  weight  of  the 
former,  N2,  is  2  X  14  =  28,  and  of  the  latter,  O2,  2  X  16  =  32. 
The  weight  of  22.4  liters  of  nitrogen  is  28  grams  and  of  oxygen 
32  grams;    the  latter  is,  accordingly,  the  heavier  gas.     Which  is 
the  heavier  gas,  carbon  dioxide,  CO2,  or  sulphur  dioxide,  862? 
The  molecular  weight  of  the  former  is  12  +  (2  X  16)  =  44,  and 
of  the  latter  32  +  (2  X  16)  =  64;  sulphur  dioxide  is  heavier. 

Pure  dry  air  is  essentially  a  mixture  of  oxygen  and  nitrogen 
molecules.  The  two  gases  are  present  in  such  proportions  that 
the  average  weight  of  the  molecules  is  28.8;  any  gas  the  molecular 
weight  of  which  is  greater  than  28.8  is  heavier  than  air. 

79.  In   commercial   work   weights   are   usually   expressed   in 
pounds  and  the  volume  of  gases  in  cubic  feet.     It  can  be  readily 
calculated  that  if  1  gram-molecular-weight  of  a  gas  has  a  volume 
of  22.4  liters,   1   pound-molecular-weight  has  a  volume  of  359 
cubic  feet.     The  latter  relation  is  commonly  made  use  of  in  the 
calculations  involved  in  technical  work. 

EXERCISES 

NOTE:  The  atomic  weight  of  the  elements  can  be  found  in  the  table  on 
the  inside  of  the  back  cover.  All  questions  in  regard  to  volume  relations 
refer  to  standard  conditions,  0°  and  760  mm. 

1.  Calculate  the  molecular  weight  and  the  percentage  of  oxygen  in  each 
of  the  following:   (a)  H2SO4,  (6)  KMnO4,     (c)  P2O6. 

2.  Calculate  the  percentage  of  each  element  in  potassium  chlorate,  KC1O3. 

3.  How  many  grams  of  hydrogen  can  be  obtained  when  25  grams  of  zinc 
are  dissolved  in  hydrochloric  acid?     Zn  +  2HC1  =  ZnCl2  +  H2. 

4.  How  much  zinc  must  be  dissolved  in  sulphuric  acid  to  produce  40  liters 
of  hydrogen? 

5.  What  weight  of  potassium  chlorate  is  required  to  furnish  enough  oxygen 
to  fill  five  250  c.c.  bottles,  assuming  that  one-quarter  of  the  gas  evolved  is 
lost  in  collecting  the  gas. 

6.  (a)  What  weight  of  carbon  is  required  to  react  with  100  pounds  of 
zinc  oxide  according   to  the  folio \ving    equation:    ZnO  +  C  =  Zn  +  CO? 
(6)  How  much  zinc  is  formed?     (c)  What  is  the  weight  of  the  carbon  mon- 
oxide formed  and  what  is  its  volume  in  cubic  feet? 


76  INORGANIC  CHEMISTRY  FOR  COLLEGES 

7.  It  was  found  by  experiment  that  1  liter  of  chlorine,  C12,  weighs  3.17 
grams.     Calculate  the  volume  of  1  gram-molecular-weight  of  the  gas. 

8.  Calculate  the  weight  of  1  liter  and  1  cubic  foot  of  each  of  the  following 
gases:    (a)  HC1,  (6)  SO2,  (c)  N2,  (d)  CO,  (e)  CO2. 

9.  Arrange  the  gases  having  the  following  formulas  in  the  order  of  increas- 
ing density:  N2,  CO,  NO,  CO2,  SO2,  NH3,  CH4. 

10.  Calculate  the  weight  of   1  liter  of  a  mixture  containing   (a)  equal 
volumes  of  oxygen  and  nitrogen,  (b)  1  volume  of  oxygen  and  4  volumes  of 
nitrogen. 

11.  (a)  Calculate  the  number  of  cubic  feet  of  air  required  to  burn  1  long 
ton  of  coke  which  contains  87  per  cent  of  carbon :  C  -f-  O2  =  CO2.     (6)  What 
is  the  relation  between  the  volume  of  the  air  used  and  the  volume  of  gases 
issuing  from  the  furnace  after  they  have  been  cooled  to  the  original  temperature 
of  the  air?     (c)  If  the  carbon  burns  to  carbon  monoxide — 2C  +  O2  =  2CO — 
what  is  the  relation  between  the  volumes  asked  for  in  (b)  above? 

12.  Calculate  the  volume  in  cubic  feet  of  1  pound-molecular-weight  if 
1  gram-molecular-weight  has  a  volume  of  22.4  liters. 

13.  An  experiment  showed  that  when  9.75  grams  of  zinc  were  dissolved  in 
acid,  0.15  gram  of  hydrogen  was  formed.     Calculate  the  atomic  weight  of 
zinc  taking  1  as  the  atomic  weight  of  hydrogen. 

14.  It  was   found  in  an    experiment  that  2  grams  of  silver  united  with 
0.657  gram  of  chlorine  to  form  silver  chloride:  2Ag  +  C12  =  2AgCl.     Assum- 
ing the  atomic  weight  of  chlorine  as  35.46,  calculate  the  atomic  weight  of 
silver. 


CHAPTER  VIII 

MEASUREMENT  OF  GASES 

80.  On  account  of  the  fact  that  gases  are  so  relatively  light 
it  is  very  difficult  to  weigh  them  with  any  degree  of  accuracy. 
The  amount  of  hydrogen  which  weighs  as  much  as  a  five-cent 
piece  has  a  volume  under  ordinary  conditions  of  over  57  liters  or 
15  gallons.     The  accuracy  of  the  results  obtained  in  weighing 
such  bulky  substances  is  markedly  affected  by  the  fact  that  the 
weight  of  the  gas  is  very  small  compared  with  the  weight  of  the 
vessel  which  contains  it.     The  greatest  experimental  skill  and  the 
most  refined  apparatus  are  necessary  to  obtain  accurate  results 
in  weighing  gases.     We  can,  however,  measure  easily  the  volume 
a  gas  occupies,  and  since  the  relations  between  the  weights  and 
volumes  of  all  gases  have  been  determined  with  great  care,  the 
weight  of  any  sample  of  gas  the  volume  of  which  is  known,  can 
be  determined  by  calculation. 

When  gases  are  heated  they  expand,  and  when  they  are  sub- 
jected to  pressure  their  volume  decreases.  For  this  reason  if  the 
weight  of  a  definite  volume  of  gas  is  to  be  recorded,  a  statement 
must  be  made  as  to  the  temperature  and  the  pressure  under  which 
the  gas  was  measured.  It  is  not  necessary,  however,  to  weigh 
gases  under  all  possible  conditions  of  temperature  and  pressure, 
because  simple  laws  have  been  discovered  in  regard  to  the  behavior 
of  gases  when  these  two  factors  change. 

81.  Thermometers. — Before  discussing  these  important  laws 
it  is  necessary  to  get  a  definite  knowledge  of  how  temperature  and 
pressure   are   measured.     Temperature   is   measured   by   a   ther- 
mometer, which  is  an  instrument  based  on  the  fact  that  certain 
substances  increase  in  volume  when  they  are  heated.     A  ther- 
mometer is  made  by  blowing  a  small  bulb  on  the  end  of  a  long 
glass  tube,  the  internal  diameter  of  which  is  very  small.     The 
bulb  and  the  lower  part  of  the  tube  are  filled  with  mercury  and 

77 


78  INORGANIC  CHEMISTRY  FOR  COLLEGES 

then  placed  in  melting  ice.  The  place  on  the  tube  where  the  sur- 
face of  the  mercury  stands  is  marked.  The  whole  apparatus  is 
next  put  in  the  steam  rising  from  boiling  water;  the  mercury 
increases  in  volume  as  it  is  heated  and  rises  in  the  tube.  When 
the  surface  of  the  mercury  no  longer  rises  the  point  where  it 
stands  is  marked.  These  two  temperatures — the  melting-point 
of  ice  and  the  boiling-point  of  water — are  the  fixed  points  on 
thermometers.  If  the  thermometer  is  to  be  graduated  with  a 
Fahrenheit  scale,  the  one  commonly  used  in  daily  life,  the  temp- 
erature of  melting  ice  is  marked  32°  and  that  of  boiling  water  212°. 
The  space  between  these  is  divided  into  180  divisions  (212 — 32) 
which  are  called  degrees.  The  stem  of  the  thermometer  is  marked 
above  the  212°  point  and  below  the  0°  point  with  divisions  equal 
in  length  to  those  between  these  two  temperatures.  It  is,  there- 
fore, possible  to  read  temperatures  with  a  mercury  thermometer 
from  the  freezing-point  to  the  boiling-point  of  the  metal.  On 
the  centigrade  thermometer  the  freezing-  and  boiling-points  of 
water  are  marked  0°  and  100°  respectively,  and  there  are  100 
degrees  between  them.  On  account  of  its  convenience  the  centi- 
grade scale  is  used  in  all  scientific  work  except  that  with  which 
the  public  come  directly  in  contact,  such  as  meteorology  (weather 
observations),  steam  engineering,  etc.  Unless  it  is  noted  to  the 
contrary  all  temperatures  recorded  in  this  book  refer  to  the  centi- 
grade scale. 

82.  Measurement  of  Pressure. — The  numerical  value  of  any 
pressure  is  usually  expressed  by  stating  the  height  of  a  column  of 
mercury  which  exerts  the  same  pressure.  When  we  say,  for 
example,  that  the  pressure  of  the  atmosphere  is  76  cm.,  we  mean 
that  the  air  presses  down  on  any  given  area  with  the  same  force 
that  a  column  of  mercury  76  cm.  high  which  covered  this  area, 
would  exert.  A  column  of  mercury  1  cm.  high  and  1  sq.  cm.  in 
area  weighs  13.5955  grams.  The  absolute  unit  of  pressure  is  that 
exerted  by  1  gram  on  1  sq.  cm.  To  express  the  pressure  of  the 
atmosphere  in  this  unit  we  must  accordingly  multiply  76  by 
13.5955;  it  is  thus  1033.3  grams  per  square  centimeter.  Since, 
however,  pressures  are  conveniently  measured  by  balancing  a  col- 
umn of  mercury  against  them,  they  are  generally  expressed  in  centi- 
meters of  mercury.  Pressures  are  expressed  in  grams  per  square 
centimeter  in  the  so-called  C.  G.  S.  system,  where  the  centimeter, 


MEASUREMENT  OF  GASES 


79 


gram,  and  second  are  the  fundamental  units.  In  expressing  pres- 
sure in  commercial  work,  such  as  the  pressure  of  steam  in  a  boiler, 
the  unit  is  the  pressure  exerted  by  one  pound  on  one  square  inch. 
When  a  larger  unit  is  required,  pressure  is  expressed  in  atmos- 
pheres; an  atmosphere  is  1033.3  grams  per  square  centimeter  or 
14.7  pounds  per  square  inch. 

83.  Atmospheric  Pressure. — All  substances  have  weight. 
The  gases  that  make  up  the  air  are  attracted  by  the  earth  and 
press  down  upon  it.  The  pressure  of  the  atmosphere  is  equal  to 
14.7  pounds  on  each  square  inch  of  the  earth's  surface;  this 
means  that  if  we  could  weigh  a  column 
of  air  one  square  inch  in  area,  which 
extended  from  the  earth  to  the  extreme 
limit  of  the  atmosphere  it  would  weigh 
14.7  pounds.  Of  course  this  cannot  be 
done  on  an  ordinary  balance,  but  it  can 
be  done  if  we  balance  the  weight  of  the 
atmosphere  against  a  column  of  mercury. 
The  instrument  used  to  do  this  is  called 
a  barometer.  In  its  simplest  form  it  is 
represented  by  Fig.  9.  It  is  made  by 
carefully  filling  with  mercury  a  long 
tube  closed  at  one  end.  The  tube  is 
then  covered  so  that  no  mercury  can 
escape,  inverted,  and  the  lower  end 
placed  under  the  surface  of  mercury 
contained  in  a  flat  dish;  the  cover  is 
next  removed.  If  the  tube  is  long 
enough  the  mercury  will  fall  until  the 
level  of  it  in  the  tube  has  reached  a  definite  height;  it  will 
then  stand  at  rest.  Since  no  air  was  allowed  to  enter  the  tube 
there  can  be  nothing  in  it  above  the  surface  of  the  mercury 
except  the  vapor  which  is  given  off  by  the  latter.  This  is  so 
very  small  in  amount  that  it  can  be  neglected  under  most 
circumstances;  the  tube  contains  in  the  upper  part  what  is  called 
a  vacuum.  What  holds  the  mercury  up  against  the  force  of 
gravity  which  tends  to  draw  it  toward  the  earth?  The  air  presses 
down  on  the  mercury  in  the  dish;  if  that  in  the  tube  fell  the  level 
of  the  mercury  in  the  dish  would  have  to  rise,  and  the  pressure  of 


FIG.  9. 


80 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


the  air  prevents  this.  Evidently  the  mercury  in  the  tube  just 
balances  the  pressure  of  the  atmosphere,  and  if  we  measure  the 
former  we  determine  as  a  result  the  latter.  If  such  an  instrument 
as  that  described  is  examined,  day  by  day,  it  will  be  found  that 
the  height  of  the  column  of  mercury  varies.  A  relationship  has 
been  discovered  between  the  height  of  the  mercury  in  a  barometer 
and  the  weather;  when  the  barometer  is  high,  that  is,  when  the 
pressure  of  the  air  is  great,  the  weather  is  apt  to  be  fair;  when  it 
falls  below  a  certain  point,  a  storm  may  be  expected.  Air  travels 
from  a  place  where  there  is  high  pressure,  that  is,  where  the  air  is 
heavy,  to  a  place  where  the  pressure  is  low  and  the  air  is  light. 
The  rushing  in  of  air  at  different  temperatures  from  all  directions 
produces  a  storm  center. 

In  order  to  express  relationships  between  the  volumes  and 
weights  of  gases,  standard  conditions  of  temperature  and  pressure 
have  been  defined;  these  are  0°  and  76  cm.  of  mercury.  The 

standard  pressure  adopted  is  the 
average  pressure  of  the  air  at  sea- 
level. 

84.  Boyle's  Law. — Robert  Boyle, 
an  English  scientist,  discovered  in 
1660  a  law  which  expresses  the  effect 
on  the  volume  of  a  definite  amount 
of  a  gas  produced  by  changing  the 
pressure  on  it.  This  effect  can  be 
demonstrated  by  a  simple  experi- 
ment. A  tube  of  the  shape  repre- 
sented in  Fig.  10  is  partly  filled  with 
mercury  at  a  as  indicated  by  the 
diagram.  It  is  closed  at  b  and 
open  at  c,  and  air  is  present  in  both 
arms  of  the  tube.  The  pressure  on 
the  surface  of  the  mercury  at  d  is 
that  of  the  atmosphere;  since  the 
surface  at  e  is  at  the  same  level  as 
FIG.  10.  FIG.  10a.  that  at  d  the  air  entrapped  in  the 

closed  end  of  the  tube  is  also  under 

a  pressure  of  one  atmosphere.  The  pressure  on  this  air  can  be 
changed  by  pouring  mercury  into  the  tube  at  the  point  marked  c. 


MEASUREMENT  OF  GASES  81 

Fig.  10a  represents  what  happens  when  this  is  done;  the  volume  of 
the  gas  decreases.  The  pressure  exerted  on  the  entrapped  gas  can 
be  determined  by  measuring  the  height  of  the  column  of  mercury 
from  /  to  g,  for  the  following  reasons :  The  pressure  down  at  h  is 
the  same  as  that  down  at  g.  In  any  liquid  the  pressures  exerted 
on  all  points  at  the  same  level  are  the  same.  The  pressure  down 
at  g  is  equal  to  that  of  the  column  of  mercury  fg  plus  that  of 
the  air  which  presses  down  at  /.  If,  in  the  experiment,  the  height 
of  fg  is  76  cm.  then  the  pressure  at  g  and  also  at  h  is  two  atmos- 
pheres, and  we  have  doubled  the  pressure  on  the  gas  in  i.  If  we 
examine  the  volume  of  the  gas  we  shall  find  that  it  is  just  one-half 
what  it  was  before.  If  a  series  of  experiments  were  carried  out 
in  which  the  gas  was  subjected  to  pressures  of  1,  2,  3,  4,  5,  etc., 
atmospheres,  it  would  be  found  that  the  volumes  were,  respect- 
ively, 1,  i,  3,  J,  £,  etc.  These  facts  can  be  stated  by  saying  that 
the  volume  of  any  sample  of  gas  varies  inversely  as  the  pressure 
on  it;  it  varies  inversely  since  as  the  pressure  increases  the  vol- 
ume decreases.  Another  way  to  express  the  fact  is  to  say  that 
the  product  of  the  volume  and  the  pressure  is  a  constant, 
1X1=1,  J  X  2  =  1,  etc.  These  general  statements  can  be 
expressed  in  the  two  forms, 

t>  cc  -     or     pv  =  c. 
P 

Where  v,  p,  and  c  represent  volume,  pressure,  and  a  constant, 
respectively.  The  symbol  oc  means  "  varies  as."  The  volume 

varies  as  -  because  the  ratio  is  an  inverse  and  not  a  direct  one. 

P 

Boyle's  law  is  usually  expressed  by  the  equation  pv  =  c.  The 
law  holds  true  only  when  there  is  no  change  in  the  temperature  of 
the  gas. 

85.  If  we  know  the  volume  of  a  gas  at  any  pressure  we  can 
calculate  its  volume  at  any  other  pressure  or  its  pressure  at  any 
other  volume.  For  any  definite  weight  of  gas  pv  remains  constant. 
Let  p  and  v  represent  the  pressure  and  volume  under  one  condi- 
tion, and  p'  and  v'  the  values  for  these  under  another  condition. 
Since  pv  =  c  and  p'v'  =  c, 

pv  =  p'v'. 


82  INORGANIC  CHEMISTRY  FOR  COLLEGES 

If  we  know  the  values  of  any  three  quantities  in  this  equation  we 
can  find  the  fourth.  The  solution  of  a  problem  will  make  this 
clear.  A  sample  of  gas  was  examined  and  found  to  have  the 
volume .  10  c.c.  when  it  was  under  a  pressure  of  75  cm.  What 
would  be  its  volume  at  76  cm.?  The  letter  p  and  v  refer  to  one 
condition  of  the  gas;  pf  and  v'  to  the  other.  Let  10  =  v,  then 
75  =  p}  and  76  =  pr.  Substitute  in  the  formula  pv  =  p'v'  and 
we  get 

75  X  10  -  76  X  x 

750  =  76x 
x  =  Iff-  =  9.86  c.c. 

Whenever  a  problem  like  this  is  solved  it  is  advisable  to 
examine  the  answer  to  determine  whether  it  is  a  reasonable  one, 
for  in  this  way  mistakes  are  often  avoided.  In  this  case  the  pres- 
sure is  increased  and,  as  a  consequence,  the  volume  should 
decrease;  the  answer  is  in  accord  with  this  and  is  probably  cor- 
rect. If  a  mistake  had  been  made  in  effecting  the  substitution 
of  numbers  for  the  letters  in  the  formula,  the  answer  might  have 
been  a  number  greater  than  10,  and  the  examination  of  the  result 
would  have  brought  out  the  fact  that  an  error  had  been  made. 

86.  Charles'  Law. — In  1787  Charles  discovered  that  the  change 
in  volume  of  a  gas  when  subjected  to  a  change  in  temperature 
could  be  expressed  by  a  simple  law.  When  the  temperature  of  a 
definite  amount  of  a  gas  is  changed  1  degree,  the  change  in  volume 
of  the  gas  is  -yfg-  of  the  volume  the  gas  would  occupy  at  0°,  pro- 
vided the  pressure  does  not  change;  if  the  temperature  is  raised 
the  gas  expands,  and  if  it  is  lowered  it  contracts.  For  example, 
if  we  have  1  c.c.  of  gas  at  0°  its  volume  at  1°  is  1  +  -^3  c.c.;  at 
2°  it  is  1  +  vfa  c.c.;  at  3°  it  is  1  +  -^  c.c.  The  volume  at 
0°,  1°,  2°,  and  3°  are  f^f,  f^,  Iff,  and  f|f,  respectively.  These 
numbers  are  evidently  in  the  ratio  273  :  274  :  275  :  276.  If  we 
change  our  scale  of  temperature  and  call  0°  centigrade  273°,  then 
1°  becomes  274°,  2°  becomes  275°,  etc.  If  we  use  such  a  scale  of 
temperature  we  see  at  once  that  the  volumes  of  the  gas  are  in  the 
same  ratio  as  their  temperatures.  The  temperature  scale  arrived 
at  in  this  way  was  found  to  be  of  great  value  in  other  connections, 
and  is  much  used  in  science;  it  is  called  the  absolute  scale.  Any 


MEASUREMENT  OF  GASES  83 

temperature  on  the  centigrade  scale  is  changed  to  the  absolute 
scale  by  adding  to  it  273°.  A  capital  letter  T  is  used  to  represent 
temperature  in  this  scale,  and  a  small  letter  I,  that  on  the  centi- 
grade scale;  T  is  thus  equal  to  t  +  273°. 

These  facts  can  be  summed  up  in  the  expression  v  oc  T  or 
v  :  v'  =  T  :  T'  where  v  is  the  volume  of  the  gas  at  the  temperature 
T  and  v'  is  its  volume  at  the  temperature  T'.  The  volumes  are 
in  the  same  ratios  as  their  absolute  temperatures. 

87.  With  the  aid  of  this  law  we  can  calculate  the  volume 
which  a  gas  will  occupy  at  any  temperature,  provided  we  know 
the  volume  it  occupies  at  any  other  temperature.     For  example, 
if  we  have  20  c.c.  of  hydrogen  at  30°  what  will  be  its  volume  at 
40°?     The  volume  at  30°  (20  c.c.)  is  to   the  volume  at  40°,  as 
30  +  273  is  to  40  +  273, 

20  :  v'  =  303  :  313 
303z/  =  6260 
v'  =  20.6  c.c. 

The  answer  should  be  inspected;  the  temperature  of  the  gas  is 
increased  and  its  volume  should,  therefore,  be  greater;  the  answer 
is  a  reasonable  one. 

88.  The  Combination  of  Boyle's  and  Charles'  Laws.— The 
mathematical  expressions  for  these  two  laws  are  as  follows: 

v  oc  -    if  the  temperature  is  constant. 
P 

v  oc  T7,  if  the  pressure  is  constant. 
These  can  be  combined  into  one  expression: 

v  oc  -  X  T,  if  pressure  and  temperature  vary.1 
pv  oc  T7, 

1  A  simple  analogy  will  make  clear  v/hy  the  volume  is  proportional  to  the 
product  of  the  terms  -  and  T.  The  areas  of  a  series  of  quadrangles  are  pro- 
portional to  their  lengths  if  their  breadths  are  constant;  the  areas  are  pro- 
portional to  their  breadths  if  their  lengths  are  constant;  and  the  areas  are 
proportional  to  the  product  of  their  lengths  and  their  breadths  if  both  these 
factors  vary. 


84  INORGANIC  CHEMISTRY  FOR  COLLEGES 

then 

pv  :  p'v'  =  T  :  T', 
or 

WL    T_ 

p'v'  ~  Tr 

This  formula  contains  six  quantities;  if  any  five  are  known  the 
sixth  can  be  determined.  It  is  generally  used  to  calculate  the 
volume  of  a  gas  under  certain  required  conditions,  when  the  volume 
is  known  under  other  conditions.  To  do  this  it  is  changed  to  the 
following  form : 


This  means  if  we  want  to  find  out  the  volume  of  a  gas  (v)  under 
certain  conditions  of  temperature  (T)  and  pressure  (p)  when  the 
volume  (t/)  at  another  temperature  (T'}  and  pressure  (p')  are 
known,  we  multiply  this  known  volume  by  two  factors,  one  of 

T_ 

r 


which  is  made  up  of  the  two  temperatures,  7FJ,  and  the  other  of 


the  two  pressures,  — .     This  is  done  in  a  very  simple  way  which 

avoids  confusion  in  the  substitution.  If  the  temperature  increases 
then  the  volume  increases,  and  T  and  T'  should  be  selected  so 

T 

that  the  value  of  the  fraction  —f  is  greater  than  1;    the  larger 

number  is  put  in  the  numerator;  and  the  reverse  is  done  if  the 
temperature  decreases.  If,  on  the  other  hand,  the  pressure  is  to 
increase,  the  volume  of  the  gas  will  be  less,  and  the  values  of  the 

pressure  are  substituted  in  the  expression  —  so  that  the  fraction 

is  less  than  1;  the  smaller  number  is  put  in  the  numerator;  if  the 
pressure  is  to  get  less  the  volume  will  increase  and  the  larger  pres- 
sure should  be  put  in  the  numerator.  A  problem  will  illustrate 
the  method.  What  will  be  the  volume  at  0°  and  76  cm.  pressure 
of  60  c.c.  of  gas  measured  at  20°  and  74  cm.  pressure?  The 
temperatures  must  be  expressed  first  in  the  absolute  scale;  they 
are  273°  (0°)  and  293°  (20°).  Since  under  the  conditions  sought 
the  temperature  is  lower,  the  volume  will  be  less,  and  T  and  T' 


MEASUREMENT  OF  GASES 


85 


are  selected  to  make  the  value  of  —,  less  than  1 ;  therefore,  we  use 
f|f.  The  pressure  is  to  increase,  and  the  volume  is,  as  a  conse- 
quence, less,  therefore  —  =  ff .  The  value  of  v  is,  accordingly, 
expressed  thus: 

v  =  GO  x  m  x  H, 


9 


v  =  54.4  c.c. 

89.  Pressure  of  Aqueous  Vapor. — Gases  are  often  measured 
in  vessels  which  are  inverted  over  water.  Under  these  circum- 
stances the  vessel  contains  not  only  the  gas,  the  volume  of  which 
is  sought,  but  water  in  the  form  of  vapor;  as  this  is  a  gas  it  exerts 
a  pressure  and  we  must  know  the  amount  of  this  if  we  are  to  know 
the  pressure  exerted  by  the  gas  itself.  That  water  gives  off  a 
vapor  which  exerts  pressure  in  the  way  a  gas  «  b 

does,  can  be  shown  by  a  simple  experiment. 
Two  barometer  tubes  are  set  up  side  by 
side  (Fig.  11).  The  level  of  the  mercury  will 
stand  at  the  same  height  in  the  two  tubes. 
A  few  drops  of  water  are  now  forced  into  tube 
b  by  placing  a  medicine-dropper  containing 
water  under  the  lower  end  of  b  in  the  mercury. 
The  water  rises  in  the  tube  through  the  mercury 
to  its  upper  surface,  and  at  once  the  level  of 
the  mercury  in  it  falls  to  c.  If  not  too  much 
water  is  added  it  will  practically  all  disappear; 
it  changes  into  a  gas  and  forces  the  mercury 
down.  The  pressure  exerted  by  the  water- 
vapor  is,  evidently,  equal  to  that  of  a  column 
of  mercury  of  the  height  equal  to  the  distance 
between  the  level  at  c  and  at  d.  If  next  we 
heat  the  tube  b  we  will  find  that  the  level 
of  the  mercury  falls;  there  will  be  a  definite 
place  at  which  the  mercury  stands  for  each  temperature.  The 
difference  between  the  height  of  the  mercury  in  the  barometer 
a  and  in  the  tube  b  represents  in  each  case  the  pressure  exerted 
by  water-vapor  at  that  temperature.  In  making  such  a  series 
of  observations  it  is  necessary  to  have  at  all  times  a  little 


FIG.  11. 


86  INORGANIC  CHEMISTRY  FOR  COLLEGES 

liquid  water  present,  so  that  as  much  vapor  as  possible  will 
form,  or,  as  it  is  generally  expressed,  the  vapor  must  be 
saturated.  If  the  tube  b  is  heated  at  100°,  the  boiling-point 
of  water,  the  level  of  the  mercury  will  be  forced  down  to  that  in 
the  dish  e;  this  means  the  pressure  of  water- vapor  is  equal  at  its 
boiling-point  to  the  pressure  of  the  air.  The  pressure  exerted  by 
water-vapor,  has  been  determined  for  all  temperatures,  and  when 
any  particular  value  is  needed  it  can  be  found  in  a  table.  (See 
Appendix.) 

90.  If  a  gas  is  measured  in  a  vessel  over  water  the  pressure 
exerted  within  the  vessel  on  the  surface  of  the  water  is  evidently 
made  up  of  the  pressure  of  the  gas  plus  the  pressure 
of  the  water-vapor.  In  order  to  simplify  the  cal- 
culations it  is  advisable  when  measuring  a  gas  over 
water  to  place  the  vessel  in  such  a  position  that 
the  level  of  the  water  inside  and  outside  the  tube  is 
the  same.  In  this  case,  Fig.  12,  the  pressure  down 
at  a  is  that  of  the  atmosphere  which  is  determined 
by  reading  a  barometer.  Since  a  and  6  are  at  the 
same  level  the  pressure  at  6  equals  that  at  a  and 
is,  accordingly,  equal  to  that  of  the  atmosphere. 
The  pressure  at  6  is  the  sum  of  two  pressures:  the 
pressure  of  the  gas  contained  in  the  tube  and  the 
pressure  due  to  the  water-vapor.  We  have  then, 


t|Y4 

=-_B- 


FIG.  12.        pressure    of   the   air  =  pressure   of  the   gas  +  the 
pressure  of  water-vapor;  the  pressure  of  the  gas  = 
pressure  of  the  air  —  pressure  of  water- vapor. 

If  the  volume  of  a  gas  measured  over  water  at  a  certain  tem- 
perature and  pressure  is  known,  it  is  possible  to  calculate  what  the 
volume  would  be  at  any  other  temperature  and  pressure  when  the 
gas  is  free  from  water.  This  is  done  by  using  the  combined 
expression  for  the  two  gas  laws  and  substituting  for  the  observed 
pressure  of  the  gas,  the  difference  between  it  and  the  pressure  of 
water-vapor  at  the  observed  temperature.  A  problem  will  be 
solved  as  an  example.  A  gas  measured  over  water  had  the  volume 
50  c.c.  at  20°  and  740  mm.  pressure.  The  pressure  of  water-vapor 
at  20°  is  17  mm.  What  would  its  volume  be  if  dry  and  at  0°  and 
760  mm.  pressure? 


MEASUREMENT  OF  GASES  87 


EXERCISES 

NOTE:  All  temperatures  given  refer  to  the  centigrade  scale  unless  a  state- 
ment is  made  to  the  contrary. 

1.  Express  in  degrees  Fahrenheit  the  following  temperatures:    (a)  30°, 
(6)  128°,    (c)  15°,    (d)   —10°.     Express  in  degrees  centigrade  the  following 
temperatures:   (e)  40°  F.,    (/)  100°  F.,    (0)   -10°  F. 

2.  If  5  liters  of  air  at  the  temperature  of  the  room,  20°,  are  cooled  to  0°, 
what  will  be  the  volume  of  the  air? 

3.  A  sample  of  gas  was  found  to  have  the  volume  250  c.c.  when  measured 
at  70  cm.  pressure.     What  would  be  its  volume   (a)  at  76  mm.  pressure,  and 
(6)  if  it  were  compressed  to  a  pressure  of  10  atmospheres? 

4.  The  observed  temperature  and  pressure  of  50  c.c.  of  a  gas  were  23° 
and  745  mm.     Calculate  the  volume  of  the  gas  at  0°  and  760  mm. 

5.  A  sample  of  a  gas  measured  over  water  was  found  to  have  the  volume 
40  c.c.  when  the  height  of  the  barometer  was  750  mm.  and  the  temperature 
was  20°.     Calculate  the  volume  of  the  dry  gas  at  0°  and  760  mm.     The 
pressure  of  aqueous  vapor  at  20°  is  17.4  mm. 

6.  If  air  at  20°  is  passed  through  a  hot  tube  and  heated  to  1000°  what 
is  the  relation  between  the  volume  of  the  gas  before  and  after  passing  through 
the  tube? 

7.  Which  is  heavier,  air  or  carbon  dioxide?      If  you  examined  the  air  in 
a  room  lighted  by  gas  would  you  expect  to  find  a  higher  percentage  of  carbon 
dioxide,  which  is  one  of  the  products  of  burning  gas,  near  the  floor  or  the  ceil- 
ing?    Give  a  reason  for  your  answer. 

8.  (a)  Calculate  the  weight  of  1  liter  of  dry  air  at  0°  and  760  mm.  pres- 
sure assuming  that  the  air  contains  79  per  cent  of  nitrogen,  N2,  20  per  cent  of 
oxygen,  C>2,  and  1  per  cent  of  argon,  A.     (6)  What  would  be  the  volume  of 
1  liter  if  it  were  heated  from  0°  to  20°?     (c)  What  would  be  the  weight  of 
1  liter  of  the  air  at  20°  and  760  mm.? 

9.  (a)  If  1  liter  of  dry  air  at  20°  and  760  mm.  is  shaken  with  enough  water 
to  saturate  it  with  water-vapor  and  is  then  collected  over  water  what  would 
be  the  volume  of  the  moist  gas  produced,  neglecting  the  solubility  of  air  in 
water?     The  pressure  of  aqueous  vapor  at  20°  is  17.4  mm.     (6)  Show  that 
dry  air  is  heavier  than  moist  air. 

10.  What  is  the  percentage  by  volume  of  the  water- vapor  in  the  moist 
air  prepared  according  to  question  9  above? 

11.  (a)  Calculate  the  volume  of  1  gram  of  steam  at  100°  and  760  mm. 
What  is  the  change  in  volume  when  1  cram  of  water  is  changed  into  steam 
at   (6)  100°  and   (c)  at  500°? 

12.  It  was  found  that  a  certain  piece  of  zinc  when  dissolved  in  hydro- 
chloric acid  furnished  70  c.c.  of  hydrogen  when  measured  over  water  at  0° 
and  750  mm.     Calculate  the  weight  of  zinc  used. 

13.  What  volume  of  hydrogen  measured  over  water  at  20°  and  750  mm. 
pressure  will  be  obtained  when  10  grams  of  zinc  dissolve  in  hydrochloric  acid? 


CHAPTER  IX 
WATER 

91.  The  significance  of  water  in  nature  was  recognized  early, 
for,  as  we  have  seen,  the  first  hypothesis  as  to  the  composition  of 
matter  was  that  all  forms  of  it  were  made  up  of  four  elements, 
earth,  air,  Ere,  and  water.     The  view  that  water  is  a  constituent 
of  matter  was  a  reasonable  one,  because  so  many  substances  which 
are  dry,  and  apparently  free  from  water,  yield  it  when  heated. 
It  was  only  after  oxygen  and  hydrogen  were  discovered  that 
water  was  shown  to  be  a  compound  of  these  two  elements.     The 
way  in  which  this  was  found  out  is  of  interest.     Priestley,  the 
discoverer  of  oxygen,  was  performing  in  1781  as  he  said  "  some 
random  experiments  to  entertain  a  few  philosophical  friends." 
He  exploded  some  hydrogen  and  oxygen  in  a  vessel  by  means  of 
an  electric  spark.     The  most  striking  effect  was  that  the  gases 
decreased  in  volume,  but  it  was  observed  that  the  sides  of  the 
vessel  were  "  bedewed  "  with  a  liquid.     No  particular  attention 
was  paid  to  this  fact,  but  Cavendish,  a  physicist  present,  was 
impressed  by  it.     He  considered  it  well  worth  future  study,  and 
later  found  that  when  two  volumes  of  hydrogen  were  exploded 
with  one  volume  of  oxygen,  the  gases  disappeared  and  the  liquid 
formed  was  water.     Lavoisier  also  made  the  same  discovery. 

92.  Occurrence  of  Water. — Water  occurs  widely  distributed 
over  the  earth's  surface,  and  as  a  gas  in  the  atmosphere.     In  the 
form  of  ice  and  snow  it  covers  lofty  mountains  and  the  surface  of 
the  earth  in  the  neighborhood  of  the  poles.     Water  is  present  in 
all  living  things;  the  fluids  of  the  body  that  are  essential  to  life- 
processes  are  chiefly  water;   the  blood,  which  carries  throughout 
the  body  the  oxygen  and  other  materials  necessary  for  life,  con- 
tains 90.3  per  cent  of  water.     The  human  body  contains  about 
65  per  cent  of  water.     Without  water,  life  as  we  know  it  would 
be  impossible.     The  substances  we  use  as  foods  contain  a  large 

88 


WATER  89 

percentage  of  water;  for  example,  there  is  present  in  potatoes  63, 
apples  65,  string  beans  90,  lettuce  95,  and  beefsteak  62  per  cent 
of  water. 

The  presence  of  such  large  masses  of  water  on  the  earth's 
surface  has  a  marked  effect  on  the  climate.  When  water  is 
heated  it  rises  more  slowly  in  temperature  than  any  other  sub- 
stance, since  it  takes  more  heat  to  raise  the  temperature  of  water 
one  degree  than  it  does  in  the  case  of  an  equal  weight  of  any  other 
substance.  As  a  consequence,  islands  and  places  near  the  sea  have 
a  more  equable  climate  than  those  regions  near  the  middle  of  a 
great  continent.  Water  and  ice  have  been  the  chief  agencies  in 
making  the  surface  of  the  earth  as  it  is  to-day.  Water  slowly 
acts  on  rocks  and  dissolves  and  carries  away  a  part  of  their  con- 
stituents; some  of  these  remain  in  the  soil  and  furnish  the  material 
which  plants  require  for  their  growth.  Great  glaciers  of  ice  have 
furrowed  the  earth's  surface,  and  produced  valleys  and  moun- 
tain ranges.  Water  confined  in  the  depths  of  the  earth  becomes 
heated  to  tremendous  temperatures,  and  finally  the  enormous 
pressures  produced  disrupt  great  masses  of  the  earth  and  as  a 
result  its  surface  has  been  greatly  changed. 

Water  is,  next  to  coal,  the  most  important  source  of  energy 
which  we  can  control  and  use  for  our  benefit.  The  energy  of 
falling  water  can  be  readily  made  available  by  allowing  it  to  turn 
a  wheel;  the  mechanical  energy  so  produced  can  be  utilized 
directly  or  converted  into  electricity,  which  can  be  distributed  to 
distant  places  for  use.  We  shall  see  later  that  one  of  the  great 
recent  developments  in  chemistry  is  the  application  of  electrical 
energy  to  the  preparation  of  many  substances  of  value.  This 
development  has  been  possible  only  through  the  utilization  of 
water-power,  and  more  and  more  the  value  of  the  energy  of  falling 
water  is  being  recognized.  When  all  the  coal  is  used  up,  in  a 
thousand  years  or  more,  the  world  must  turn  to  the  waterfall 
for  its  energy,  unless  some  new  force  is  discovered.  There  are 
great  stores  of  energy  in  the  sunlight,  but  no  one  has  yet  discovered 
how  to  use  this  energy  advantangeously. 

93.  Physical  Properties  of  Water. — Water  is  a  liquid  which 
has  a  bluish-green  color  when  examined  in  thick  layers.  It  is 
taken  as  the  substance  upon  which  to  base  standards  for  measur- 
ing a  number  of  physical  properties.  Its  freezing-point  and  boil- 


90  INORGANIC  CHEMISTRY  FOR  COLLEGES 

ing-point  at  760  mm.  pressure  are  defined  as  0°  and  100°,  respect- 
ively. The  weight  of  1  c.c.  of  water  at  4°  was  the  original  de- 
finition of  a  gram,  but  later  the  weight  of  a  definite  piece  of 
platinum  based  on  the  unit,  which  is  kept  in  Paris,  was  taken  as 
the  standard.  The  calorie  is  defined  as  the  amount  of  heat  re- 
quired to  raise  the  temperature  of  1  gram  of  water  1  degree,  from 
15°  to  16°;  since  specific  heat  is  defined  as  the  quantity  of  heat 
required  to  raise  the  temperature  of  1  gram  of  a  substance  1 
degree,  the  specific  heat  of  water  is,  accordingly,  1  at  15°.  Water 
is  the  reference  substance  used  in  expressing  specific  gravity 
(176) ;  the  value  for  the  specific  gravity  for  water  is,  therefore,  1. 

As  in  the  case  of  other  liquids  the  density  (175)  of  water 
changes  with  temperature,  but  the  change  in  this  case  is  not  uni- 
form; since  the  density  of  a  liquid  is  defined  as  the  weight  of  1  c.c. 
of  it,  the  density  of  water  is  1  at  4°;  it  decreases  as  the  tempera- 
ture rises  or  falls  from  this  point;  at  0°  it  is  0.99987  and  at  100° 
it  is  0.95838. 

When  water  freezes  it  expands;  1  gram  of  ice  at  0°  has  the 
volume  1.0908  c.c.  The  heat  of  fusion  of  1  gram  of  ice  is  79 
calories,  and  the  heat  of  vaporization  of  1  gram  of  water  at  its 
boiling-point  is  540  calories. 

94.  Water- Vapor. — The  fact  has  been  brought  out  that  when 
water  evaporates  it  passes  into  the  air  as  a  gas.     This  phenomenon 
has  a  marked  effect  on  the  climate.     Through  the  evaporation  of 
the  water  on  the  earth's  surface  it  passes  into  the  air,  and  later, 
as  the  result  of  condensation,  forms  clouds  and  finally  returns  as 
rain  or  snow.     The  formation  of  dew  when  the  air  cools  as  the 
sun  sinks  and  the  heat  reaching  the  earth  lessens,  is  the  result  of 
the  change  of  the  water-vapor  to  a  liquid.     In  this  way  growing 
plants  get  much  of  the  water  they  require  when  it  is  not  furnished 
as  rain.     Equally  important  is  the  fact  that  plants  take  up  directly 
from  the  air  water- vapor  and  carbon  dioxide,  and  transform  them 
into    the    organic   materials  of  which    they  are   made,   such  as 
wood,  starch,  etc. 

95.  The  body  contains  much  water,  and  that  at  the  surface  is 
continually  evaporating  and  passing  into  the  air.     Our  comfort 
is  materially  affected  by  the  amount  of  water-vapor  around  us. 
If  the  air  is  saturated,  that  is,  contains  as  much  as  it  can  hold  in 
the  form  of  gas,  this  evaporation  cannot  take  place  and  we  are 


WATER  91 

uncomfortable  as  a  result;  we  dread  a  "  muggy  "  day.  In  a  dry 
climate  the  water  evaporates  rapidly  and  in  so  doing  absorbs  heat 
from  the  body;  as  a  consequence,  we  do  not  feel  the  effect  of  a  high 
temperature  as  we  do  when  the  humidity  is  great.  Air  usually 
contains  about  two-thirds  of  the  maximum  amount  of  water  it 
can  hold  as  vapor;  under  these  circumstances  the  relative  humidity 
is  said  to  be  66  per  cent  (two-thirds).  The  rate  at  which  water 
evaporates  is  determined  not  by  the  absolute  amount  of  water 
present  in  the  air,  but  by  the  relation  between  the  amount  present 
and  the  amount  the  air  can  hold  at  the  temperature  which  pre- 
vails. Since  our  comfort  is  determined  by  the  rate  at  which 
water  evaporates  from  the  skin,  the  relative  humidity  is  the  impor- 
tant factor.  When  the  relative  humidity  drops  to  40  per  cent  the 
air  feels  dry,  and  at  about  80  per  cent  it  appears  to  be  damp. 

96.  The  amount  of  water  which  can  exist  as  vapor  in  the  air  is 
determined  by  the  temperature.     This  can  be  demonstrated  by 
the  use  of  the  experiment  already  described  in  connection  with 
the  determination  of  the  pressure  of  water-vapor  (89).     A  small 
amount  of  water  is  allowed  to  rise  through  mercury  in  a  barometer 
tube  and  when  it  reaches  the  surface  a  part  evaporates  and  the 
vapor  produced  depresses  the  mercury  as  the  result  of  the  pressure 
it  exerts.     As  the  temperature  is  raised,  more  and  more  water 
evaporates  and  the  level  of  the  mercury  sinks;  if  the  temperature 
is  lowered  water  condenses  and  the  mercury  rises.     The  fact  that 
the  maximum  amount  of  water- vapor  held  in  air  varies  with  the 
temperature,  is  the  cause  of  important  natural  phenomena.     A 
cloud  can  form  in  a  blue  sky  as  the  result  of  the  lowering  of  tem- 
perature produced  by  a  current  of  cooler  air.     The  formation  of 
dew  has  been  referred  to  and  we  see  now  that  it  is  the  result  of  the 
drop  in  temperature  which  comes  with  the  setting  sun.     In  sum- 
mer, water  soon  collects  on  a  vessel  containing  ice- water;    this 
occurs  when  the  temperature  of  the  air  next  the  vessel  is  reduced 
to  that  at  which  the  air  is  saturated  by  the  water  present  in  it. 
As  the  temperature  falls  below  this  point  the  excess  water  precipi- 
tates out  as  drops.     The  temperature  at  which  moisture  is  first 
visible  on  a  smooth  surface  when  air  is  cooled  in  contact  with  it, 
is  said  to  be  the  dew-point  of  the  air. 

97.  Chemical  Conduct  of  Water. — Water  is  a  very  stable  sub- 
stance.   When  it  is  formed  from  hydrogen  and  oxygen  a  very 


92  INORGANIC  CHEMISTRY  FOR  COLLEGES 

large  amount  of  chemical  energy  is  lost  as  heat.  In  order  to 
decompose  water  it  must  be  heated  to  a  very  high  temperature 
or  brought  into  contact  with  active  substances  possessing  a  large 
amount  of  chemical  energy.  When  water  is  heated  to  2000°  it 
dissociates  into  hydrogen  and  oxygen  to  the  extent  of  less  than 
2  per  cent. 

Water  reacts  with  certain  metals  in  the  way  already  explained 
in  section  43.  In  these  cases  the  oxygen  unites  with  the  metal 
and  hydrogen  is  set  free;  the  equation  for  the  reaction  with  zinc 
is  as  follows: 

Zn  +  H2O  =  ZnO  +  H2 

Water  combines  with  oxides  and  forms  bases  and  acids.  Quick- 
lime, for  example,  is  the  oxide  of  calcium;  it  is  converted  by 
water  into  slaked  lime,  calcium  hydroxide,  which  is  a  base: 

CaO  +  H2O  =  Ca(OH)2 

Sulphur  dioxide,  formed  by  burning  sulphur  in  air,  unites  with 
water  to  form  sulphurous  acid: 

S02  +  H20  =  H2S03 

It  will  be  recalled  that  it  was  reactions  of  the  latter  class  that 
suggested  to  Lavoisier  the  name  for  oxygen;  he  found  that  when 
certain  substances  burned  in  the  gas  the  products  formed  reacted 
with  water  to  form  acids;  oxygen  was,  thus,  the  acid-former.  It 
was  found  later,  however,  that  many  acids  do  not  contain  oxygen. 
98.  Hydrates. — Water  unites  directly  with  certain  substances 
and  forms  compounds  which  are  more  or  less  readily  broken  down, 
by  heating  at  moderate  temperatures,  into  water  and  the  original 
salts.  Copper  sulphate,  CuSC>4,  is  a  white  powder,  which  dis- 
solves in  water  and  forms  a  blue  solution.  When  the  latter 
evaporates  slowly  blue  crystals  are  formed,  which  have  the  com- 
position represented  by  the  formula  CuSOi,  5H2O.  If  these  crys- 
tals are  heated  at  a  temperature  somewhat  above  the  boiling- 
point  of  water,  they  crumble,  water  is  lost,  and  the  anhydrous 
salt  is  obtained.  Other  salts  behave  in  a  similar  way.  As  com- 
pounds which  contain  water  separate  from  their  solutions  as 
crystals,  the  water  in  combination  was  formerly  called  water  of 


WATER  93 

crystallization.  The  formation  of  crystals  is  not  dependent  upon 
the  presence  of  water,  however.  Anhydrous  copper  sulphate,  for 
example,  can  be  obtained  as  white,  needle-like  crystals.  For  this 
reason  we  now  speak  of  water  of  hydration,  rather  than  water 
of  crystallization.  We  have  anhydrous  copper  sulphate,  and 
hydrated  copper  sulphate.  To  express  the  fact  that  the  salt  has 
the  formula  CuSC>4,  5H2O  we  call  it  a  pentahydrate ;  a  salt  con- 
taining one  molecule  of  water  of  hydration  is  a  monohydrate,  one 
with  two  is  a  dihydrate,  with  three  a  trihydrate,  etc. 

The  pentahydrate  of  copper  sulphate  is  a  definite  chemical 
compound;  water  is  not  present  in  it  as  such.  It  contains  hydro- 
gen and  oxygen  in  the  proportion  represented  by  its  formula,  but 
we  do  not  know  how  the  atoms  are  united  to  one  another  or  to 
the  other  atoms  present.  We  write  the  formula  as  we  do  to 
express  the  fact  that  when  the  compound  is  heated  it  readily 
breaks  down  into  the  compounds  CuSO4  and  H^O. 

99.  Certain    crystalline    substances    contain   water   mechani- 
cally held  within  the  crystal.     When  a  piece  of  rock-salt  is  heated, 
water  is  given  off.     In  this  case  the  water  passes  into  steam  and 
sufficient  pressure  is  produced  to  cause  slight  explosions  which 
shatter  the  large  crystals — the  substance  is  said  to  decrepitate. 
There  is  no  definite  relation  between  the  weight  of  the  water 
included  mechanically  within  the  crystal  and  the  weight  of  the 
salt.     With  hydrated  salts  we  can  always  express  the  relation 
between  the  weight  of  the  water  of  hydration  and  that  of  the 
anhydrous  compound  by  a  definite  chemical  formula  which  has  a 
quantitative  significance.     Most  hydrated  salts  lose  their  water 
of  hydration  when  heated  to  100°. 

100.  Efflorescence. — If  a  crystal  of  washing  soda  is  left  in  the 
air  for  some  time  it  loses  its  crystalline  form  and  changes  to  a 
white  powder.     This  is  due  to  the  fact  that  at  room  temperature 
it  loses  a  part  of  its  water  of  hydration;  the  salt  having  the  formula 
Na2CO3,  10H2O  dissociates  into  water  and  the  monohydrate  of 
sodium  carbonate,  Na2CO,3,  EbO.     When  hydrated  copper  sul- 
phate, commonly  called  blue  vitriol,  CuSCU,  5H2O,  is  left  in  the 
air,  the  crystals  do  not  lose  their  form.     We  express  these  facts 
by  saying  that  washing  soda  effloresces. 

We  have  learned  that  water  gives  off  a  vapor  which  passes 
into  the  air.  Hydrated  salts  decompose  so  readily  into  water 


94  INORGANIC  CHEMISTRY  FOR  COLLEGES 

and  the  anhydrous  salt  that  at  room  temperature  they  give  off 
water-vapor,  the  pressure  of  which  can  be  measured.  This  can 
be  done  in  the  same  way  that  the  pressure  of  water-vapor  was 
determined  (89)  by  placing  a  crystal  of  the  substance  in  a  barome- 
ter tube  and  noting  the  fall  in  the  level  of  the  mercury.  If  the 
pressure  of  the  water-vapor  in  the  air  is  less  than  the  pressure  of 
the  water-vapor  from  the  hydrated  salt,  the  latter  will  decompose 
and  water  will  pass  into  the  air.  Under  ordinary  conditions 
air  is  about  two-thirds  saturated  with  water-vapor.  At  20° 
(68°  F.)  the  pressure  of  water-vapor  in  the  air  is  in  the  neigh- 
borhood of  12  mm.  At  this  temperature  the  pentahydrate  of 
copper  sulphate,  CuSCU,  SEbO,  exerts  a  vapor  pressure  of 
5  mm.;  it  does  not  effloresce.  The  decahydrate  of  sodium  sul- 
phate, Na2SO4,  lOEbO,  exerts  a  vapor  pressure  of  14  mm.  It  is 
seen  from  the  above  that  a  salt  may  effloresce  on  one  day  and  not 
on  another. 

101.  An  application  of  hydrated  salts  is  made  in  the  chemical 
laboratory  in  drying  gases.     Anhydrous  calcium  chloride  unites 
with  water  to  form  the  hexahydrate  CaCL?,  GH^O.     When  air 
containing  moisture  comes  in  contact  with  the  anhydrous  chloride 
the  water  is  absorbed  and  the  pressure  of  its  vapor  reduced  to 
that  of  the  hydrated  chloride;  as  this  is  exceedingly  small,  the  air 
is  deprived  of  nearly  all  of  its  water.     Certain  salts  take  up  so 
much  water  from  moist  air  that  they  dissolve  in  the  water;  such 
salts  are  said  to  be  deliquescent. 

102.  Composition  of  Water. — The  proportion  of  hydrogen  and 
oxygen  in  water  can  be  determined  by  analysis  or  synthesis.     In 
the  first  way  the  elements  are  separated  in  the  free  condition 
from  water  or  are  converted  into  other  substances  the  composition 
of  which  is  known;   in  the  second,  water  is  prepared  from  known 
amounts  of  other  substances.     The  composition  of  water  has  been 
determined  in  many  ways  and  with  the  greatest  accuracy.     A 
few  will  be  sketched  briefly. 

An  experiment  has  been  described  to  illustrate  the  fact  that  when 
hydrogen  reduces  copper  oxide,  water  is  formed  (55).  The  com- 
position of  water  has  been  determined  by  a  quantitative  study  of 
this  reaction,  which  is  represented  by  the  equation, 

CuO+H2  =  Cu  +  H20. 


WATER 


95 


This  was  done  in  the  following  way:  The  tube  containing  the 
copper  oxide  was  weighed  before  and  after  the  reaction  took  place 
(Fig.  13) ;  the  loss  in  weight  was  evidently  equal  to  the  weight  of 
the  oxygen  changed  into  water.  The  amount  of  the  water  was 
determined  by  absorbing  it  in  a  tube  represented  at  a  in  the 
diagram,  which  was  filled  with  phosphorus  pentoxide.  This  tube 
was  weighed  before  and  after  the  reaction,  and  the  increase  in 
weight  was,  evidently,  the  weight  of  the  water  formed.  In  this 
way  the  weight  of  the  oxygen  and  the  weight  of  the  water  were 
determined;  the  difference  between  these  two  was  the  weight  of 
the  hydrogen. 


FIG.  13. 

A  synthesis  of  water  has  been  carefully  made  in  which  all 
three  substances  involved  were  weighed — the  hydrogen,  oxygen, 
and  the  water.  The  results  of  the  most  carefully  carried  out 
experiments  lead  to  the  conclusion  that  hydrogen  and  oxygen 
unite  to  form  water  in  the  proportion  by  weight  of  2  of  the  former 
to  15.879  of  the  latter,  or  2.016  to  16. 

103.  The  composition  of  water  has  also  been  determined  by 
volumetric  analysis,  that  is,  the  volume  relations  have  been 
studied.  It  has  been  found  that  2.0027  volumes  of  hydrogen 
unite  with  1  volume  of  oxygen  to  form  water.  If  it  is  desired  to 
compare  the  relation  between  the  volumes  of  the  gases  which 
interact  and  the  volume  of  the  water  formed,  the  experiment 
should  be  carried  out  at  such  a  temperature  that  the  water  is  a 
gas — at  100°  or  above.  Under  these  conditions  it  has  been  found 
that  the  volumes  are  almost  exactly  in  the  ratio  of  two  of  hydro- 
gen, one  of  oxygen,  and  two  of  steam.  We  shall  find  that  this 
fact  becomes  of  the  greatest  importance  when  the  method  of 
determining  atomic  weights  is  considered. 


96  INORGANIC  CHEMISTRY  FOR  COLLEGES 

104.  Natural  Waters. — Chemically  pure  water  is  very  diffi- 
cult to  obtain.     Natural  waters  which  have  come  into  contact 
with  the  earth  contain  varying  amounts  of  substances  which  are 
dissolved  from  the  soil  and  the  air.     Rain-water  is  perhaps  the 
purest  natural  water,  but  it  contains  dust,  bacteria,  ammonium 
compounds,  and  other  substances  which  are  present  in  the  air.    If 
it  is  collected  after  rain  has  fallen  for  some  time  it  is  practically  free 
from  these  substances  and  is  the  purest  water  that  can  be  obtained 
without  special  precautions;    it  contains,   of  course,   the  gases 
oxygen,  nitrogen,  and  carbon  dioxide  which  are  present  in  the  air. 

The  materials  dissolved  from  the  earth  find  their  way  into  the 
ocean,  which  contains  about  3.6  per  cent  of  substances  which  are 
solids.  About  2.7  per  cent  of  sea-water  is  common  salt,  the  rest 
of  the  dissolved  solids  being  chiefly  chlorides  and  sulphates  of 
magnesium  and  calcium.  The  Great  Salt  Lake  of  Utah  contains 
23  per  cent  of  solid  matter  in  solution.  The  water  of  lakes  and 
rivers  which  are  used  directly  as  sources  of  supply  varies  in  the 
amount  of  solid  matter  present.  In  places  where  the  chief  rocks 
are  sandstone  or  granite,  the  water  contains  less  dissolved  material 
than  that  in  limestone  regions.  We  shall  see  later  that  water 
which  contains  carbon  dioxide  dissolves  limestone  slowly,  and  the 
soluble  calcium  salts  formed  render  the  water  hard. 

105.  Water  is  freed  from  impurities  by  distillation  (181).   When 
it  is  heated  the  air  dissolved  in  it  first  separates  in  bubbles  and 
then  escapes  as  the  temperature  is  raised.     At  100°  it  changes 
into  steam,  which  is  condensed  by  surrounding  with  cold  water 
the  tube  through  which  it  passes  (Fig.  14).     The  solid  materials 
are  not  volatile  at  100°  and  remain  in  the  flask  from  which  the 
water  was  distilled.     Water  treated  in  this  way  may  be  considered 
pure  enough  for  most  purposes;    it  is,  however,  far  from  abso- 
lutely pure.     To  remove  the  last  traces  of  foreign  substances  the 
water  should  be  heated  with  a  powerful  oxidizing  agent  to  destroy 
certain  organic  substances  present,  and  be  condensed  in  a  vessel 
made  of  tin  or  preferably  of  platinum,  because  water  dissolves 
traces  of  solid  substances  from  glass.     When  left  in  contact  with 
air,  water  slowly  absorbs  it.    The  taste  of  water  is  largely  due  to 
the  air  it  contains;   freshly  distilled  water  tastes  flat,  and  before 
being  used  for  drinking  purposes  it  is  aerated  by  allowing  it  to 
flow  over  charcoal  in  the  presence  of  air. 


WATER 


97 


106.  It  is  often  necessary  to  use  as  a  water  supply  a  river  into 
which  sewage  has  been  deposited.  If  the  place  of  pollution  is  far 
enough  removed,  the  water  may  be  entirely  safe  for  domestic  use, 
for  changes  take  place  in  flowing  water  which  result  in  the  con- 
version of  sewage  into  simple  and  harmless  inorganic  compounds. 
River-water  usually  contains  suspended  matter,  which  makes  it 
more  or  less  muddy  and  undesirable  for  household  use.  Many 
cities  which  are  compelled  to  draw  their  water  from  rivers,  subject 
it  to  an  elaborate  method  of  purification  to  remove  suspended 


FIG.  14. 


matter  and  any  harmful  contamination  produced  by  sewage,  which 
often  introduces  disease  germs  into  water.  Cholera  and  typhoid 
fever  are  spread  through  the  use  of  drinking  water  which  has 
been  infected  by  people  suffering  from  these  diseases. 


EXERCISES 

1 .  Why  does  moisture  collect  on  a  pitcher  containing  ice-water  in  summer 
but  not  in  winter? 

2.  Calculate  the  weight  of  1  c.c.  of  ice  at  0°  if  1  gram  has  the  volume 
1.0908  c.c. 

3.  Calculate  the  weight  in  pounds  of  1  cubic  foot  of  ice. 


98  INORGANIC  CHEMISTRY  FOR  COLLEGES 

4.  What  change  in  volume  takes  place  when  100  c.c.  of  water  at  4°  are 
changed  into  ice  at  0°? 

5.  When  a  body  floats  in  water  the  weight  of  the  body  equals  the  weight 
of  the  water  that  has  the  same  volume  as  that  part  of  the  body  which  is  under 
the  water,     (a)  Taking  the  density  of  water  as  1  and  ice  as  0.9  at  0°,  what 
would  be  the  volume  of  the  ice  above  the  surface  if  a  piece  weighing  1000 
grams  was  floated  on  water  at  0°?     (6)  If  the  block  of  ice  is  in  the  form  of 
a  cube,  how  thick  would  be  the  layer  of  ice  above  the  surface  of  the  water? 
(c)  What  percentage  of  the  ice  is  above  the  water?     (d)  What  weight  must 
be  placed  on  the  ice  to  just  cause  it  to  sink?     (e)  What  weight  of  floating 
ice  is  necessary  to  just  support  a  man  weighing  150  pounds?     (/)  If  this  is 
in  the  form  of  a  cube  what  is  its  size? 

6.  Why  is  frost  more  likely  to  be  formed  on  a  cold  night  when  the  air 
during  the  day  was  comparatively  dry  than  when  it  contained  a  large  amount 
of  water-vapor? 

7.  (a)  Calculate  the  percentage  of  hydrogen  and  oxygen  in  water  from 
the  following  results  which  were  obtained  in  an  experiment  like  that  outlined 
in  section  102.     Weight  of  copper  oxide  used  50.250  grams,  weight  after 
hydrogen  had  passed  over  it  48.500  grams,  weight  of  water  formed  1.970 
grams,     (6)  Calculate  the  percentage  composition  of  water  from  its  formula. 


CHAPTER  X 
CHLORINE.    VALENCE 

107.  Chlorine  is  an  interesting  and  important  substance;  it  is 
one  of  the  most  active  elements,  and  enters  into  reactions  with  a 
great  variety  of  other  substances;  it  is  present  in  a  large  number 
of  compounds.     A  study  of  chlorine  and  the  simpler  substances 
containing  it  will  introduce  us  to  a  number  of  typical  chemical 
changes,  and  we  shall  gain  thereby  a  deeper  knowledge  of  how 
matter  undergoes  transformation  under  the  influence  of  the  chemi- 
cal energy  which  it  contains  and  the  outside  energy  brought  to 
bear  upon  it.     But  there  is  a  practical  aspect  also  to  the  study  of 
chlorine;  the  element  and  a  number  of  its  compounds  find  impor- 
tant applications  which  are  of  great  service  to  man.    It  is  highly 
probable  that  new  uses  will  be  found  for  this  very  active  element, 
for  activity  can  be  made  useful. 

108.  Occurrence  of  Chlorine. — Chlorine  is  a  heavy  greenish- 
yellow  gas  with  a  stifling  odor.     It  does  not  occur  naturally  in  the 
free  condition;  for  if  the  element  were  set  free  in  the  air  it  would 
soon  find  something  with  which  to  unite.     Chlorine  forms  com- 
pounds with  metals  which  are  called  chlorides;  a  number  of  these 
occur  in  nature,  and  some  are  widely  distributed.     Sodium  chlo- 
ride, NaCl,  which  is  common  salt,  occurs  in  large  quantities  in 
sea-water,  of  which  it  forms  2.7  per  cent;    the  total  amount  of 
solids  present  is  3.6  per  cent.     Calcium  chloride,  CaCl2,  mag- 
nesium chloride,  MgCk,  and  potassium  chloride,  KC1,  are  also 
found  in  the  ocean.     These  chlorides  and  other  salts  occur  in  salt 
deposits,  which  have  been  produced,  in  all  probability,  as  the 
result  of  the  drying  up  of  inland  seas.     Silver  chloride,  AgCl,  is 
an  important  source  of  silver. 

On  account  of  its  wide  occurrence  and  its  abundance,  sodium 
chloride  is  the  cheapest  compound  containing  chlorine;  it  is,  as  a 
consequence,  the  substance  from  which  chlorine  is  obtained  on 

99 


100  INORGANIC  CHEMISTRY  FOR  COLLEGES 

the  commercial  scale,  either  directly  or  indirectly.  This  fact  is 
expressed  briefly  by  saying  that  sodium  chloride  is  the  source  of 
chlorine. 

109.  History   of   Chlorine. — Scheele,    a   Swedish   apothecary, 
first  separated  chlorine  from  one  of  its  compounds  in  1774.     He 
was  studying  a  mineral  called  pyrolusite  and  found  that  when  it 
was  treated  with  muriatic  acid,  a  yellow  gas  with  a  stifling  odor 
was  formed.     Pyrolusite  was  shown  later  to  be  a  dioxide  of  an 
element  called  manganese,  and  to  have  the  formula  Mn02.     Muri- 
atic acid  was  so  called  because  it  was  obtained  from  salt,  which  is 
found  in  the  sea  (Latin  murias)',   it  is  what  is  now  called  hydro- 
chloric acid  and  has  the  formula  HC1.     The  new  substance  was 
named  chlorine   (from  x^°P^,  greenish-yellow).      On  account  of 
the  fact  that  it  was  prepared  with  the  aid  of  a  substance  containing 
oxygen  it  was  thought  to  contain  this  element,  but  after  a  very 
thorough  study  of  chlorine  by  Sir  Humphry  Davy  it  was  recog- 
nized as  an  element.1 

110.  Preparation    of    Chlorine*     (a)  Electrolysis    of    Hydro- 
chloric Acid. — Chlorine  cannot  be  conveniently  prepared  by  heat- 
ing a  chloride;    mercuric  oxide  yields  oxygen  when  heated,  but 
chlorine  cannot  be  prepared  by  a  similar  decomposition  of  mer- 
curic chloride.     We  accordingly  use  another  form  of  energy  to 
effect  the  decomposition.     When  electricity  is  passed  through  a 
solution  of  hydrochloric  acid,  which  is  a  compound  of  hydrogen 
and  chlorine,  the  acid  is  decomposed  into  its  constituents.     The 
experiment  can  be  carried  out  in  a  Hofmann  apparatus,  but  the 
platinum  electrodes  should  be  replaced  by  ones  of  carbon,  since 
chlorine  unites  directly  with  the  metal  to  form  platinum  chloride, 
which  is  soluble  in  water,  and,  as  a  consequence,  the  valuable 
metal  is  slowly  dissolved.     When  the  experiment  is  carried  out 
and  precautions  are  taken  to  eliminate  a  disturbing  factor  intro- 
duced through  the  fact  that  chlorine  is  much  more  soluble  in 
water  than  hydrogen,  it  will  be  found  that  as  a  result  of  the  decom- 
position the  two  gases  are  formed  in  equal  volumes.     This  method 
of  making  chlorine  recalls  one  of  the  methods  of  preparing  oxygen 
from  water.     In  either  case  the  compound  of  the  element  with 
hydrogen  was  decomposed  into  its  constituents  by  electricity. 

1  Davy's  description  of  his  experiments  can  be  found  in  Volume  9  of  the 
Alembic  Club  Reprints.  It  is  of  interest  to  read  by  what  experimental  and 
logical  processes  a  substance  was  shown  to  be  an  element. 


CHLORINE.     VALENCE  101 

Chlorine  is  prepared  on  the  commercial  scale  by  the  electrolysis 
of  solutions  of  sodium  chloride;  as  a  result,  in  addition  to  chlorine 
are  obtained  sodium  hydroxide  and  hydrogen,  both  of  which  are 
valuable  products.  The  details  of  fe :  process  avej,  given  in  sec- 
tion 605. 

111.  (6)  Deacon  Process. — Chjo#n$  enrobe'  ser?a1^t6d.from  the 
hydrogen  with  which  it  is  united  in  hydrochloric  acid,  through 
the  agency  of  chemical  energy.     When  a  mixture  of  air  and  hydro- 
chloric acid  heated  to  about  400°  is  passed  through  a  tube  filled 
with  clay  balls  which  have  been  soaked  in  a  solution  of  copper 
chloride  and  then  dried,  chlorine  is  formed  as  the  result  of  a  reac- 
tion expressed  by  the  following  equation: 

4HC1  +  O2  =  2H2O  +  2C12 

The  finely  divided  copper  chloride  deposited  on  the  balls  of  clay 
serves  as  a  catalytic  agent.  Chlorine  is  prepared  in  this  way  by 
the  oxidation  of  hydrochloric  acid;  oxygen  is  the  oxidizing  agent, 
and  the  products  of  oxidation  are  water  and  chlorine. 

Chlorine  made  by  this  process  is  mixed  with  the  nitrogen  gas 
present  in  the  air  which  furnished  the  oxygen;  the  amount  of 
nitrogen  is  great,  since  four-fifths  of  any  volume  of  air  is  this  gas. 
This  way  of  making  chlorine  is  known  as  the  Deacon  process;  it 
was  formerly  much  used  in  England  to  prepare  chlorine  for  making 
bleaching  powder,  but  other  methods  yield  chlorine  in  a  purer 
condition,  and  are  now  used  almost  exclusively. 

112.  (c)  Manganese  Dioxide  and  Hydrochloric  Acid. — It  has 
been  stated  that  no  chlorides  are  available  which  yield  chlorine 
readily  on  heating.     Chlorides  can  be  prepared,  however,  which 
are  so  unstable  that  they  undergo  decomposition  spontaneously 
when  they  are  formed.     Reactions  of  this  type  furnish  conven- 
ient methods  of  preparing  chlorine.     One  of  these  reactions  is  the 
one  which  led  to  the  discovery  of  the  gas.     When  manganese 
dioxide  is  treated  with  hydrochloric  acid,   chlorine,   manganese 
chloride,  and  water  are  formed.     It  is  believed  that  the  reaction 
takes  place  in  two  stages;  manganese  tetrachloride  is  first  formed, 
and  then  breaks  down  into  manganese  dichloride  and  chlorine: 

MnO2  +  4HC1  =  MnCU  +  2H2O 
MnCU  =  MnCl2  +  C12 


102  INORGANIC  CHEMISTRY  FOR  COLLEGES 

If  the  reaction  is  carried  out  at  a  very  low  temperature  manganese 
tetrachloride  is  formed;  if  now,  the  temperature  is  allowed  to  rise 
the  tetrachloride  loses  chlorine  in  the  way  indicated  by  the  last 
equation  ab£>ve.\  A  single;;  equation  can  be  written  to  express  the 
two  reactions.  Sinc'e  the  'man'ganese  tetrachloride  decomposes  as 
soon  as  formed,;  ife-fbrmijlja  should  nbt  appear  in  the  equation  which 
is  to  represent  trie  'formulas  o'f  the  substances  used  and  those 
obtained.  The  equation  is,  accordingly, 

MnO2  +  4HC1  =  MnCl2  +  C12  +  2H2O 

Certain  other  dioxides,  such  as  lead  dioxide,  behave  in  a  similar 
way  when  treated  with  hydrochloric  acid. 

113.  Large  quantities  of  chlorine  are  manufactured  in  Eng- 
land by  the  reactions  indicated  by  the  equation  given  above. 
Manganese  dioxide  occurs  abundantly  in  nature  and  is  conse- 
quently cheap  in  price,  and  hydrochloric  acid  is  a  by-product 
which  is  produced  in  large  quantities  in  the  industrial  method  of 
manufacturing  soda  used  in  England.     Soda  is  not  made  in  this 
way  in  America,  and,  as  a  consequence,  hydrochloric  acid  is  not 
available  at  a  low  price.      Water-power  to  make  electricity  is 
available,  however,  and,  as  a  result,  chlorine  is  manufactured  in 
America  almost  exclusively  by  the  electrolysis  of  sodium  chloride. 
This   process   yields   a   valuable   by-product,  sodium  hydroxide, 
NaOH,  which  is  commonly  called  caustic  soda  in  trade.     It  is 
evident  that  the  particular  processes  used  to  prepare  chemical 
substances  on  the  large  scale  for  commercial  purposes  are  deter- 
mined in  the  last  analysis  by  cost,  in  which  must  be  taken  into 
account  such  factors  as  source  of  materials,  expense  of  transporta- 
tion and  labor,  availability  of  energy  derived  from  heat  and  elec- 
tricity, value  of  by-products  produced,  etc.     The  preparation  of 
useful  substances  from  this  economic  point  of  view  is  considered 
in  industrial  chemistry. 

114.  Chlorine  is  prepared  in  the  laboratory  from  hydrochloric 
acid  and  manganese  dioxide,  or  by  a  slight  modification  of  this 
method.     Instead  of  using  hydrochloric  acid  which  has  been  pre- 
pared previously,  this  compound  is  made  in  a  vessel  containing 
manganese  dioxide.     The  preparation  of  hydrochloric  acid  will  be 
discussed  in  detail  later;  it  is  sufficient  to  note  here  that  the  com- 
pound is  formed  when  sodium  chloride,  which  is  common  salt,  is 


CHLORINE.     VALENCE  103 

treated  with  sulphuric  acid.  The  reaction  that  takes  place,  in 
which  sodium  sulphate  and  hydrochloric  acid  are  formed,  is  repre- 
sented by  the  following  equation: 

2NaCl  +  H2SO4  =  2HC1  +  Na2SO4 

It  should  be  observed  that  the  sodium  and  hydrogen  atoms  change 
places.  The  hydrochloric  acid  produced  in  this  way  reacts  with 
the  manganese  dioxide  present  according  to  the  equation  already 
given : 

MnO2  +  4HC1  =  MnCl2  +  C12  +  2H2O 

This  reaction  leads  to  the  formation  of  manganese  chloride. 
Since  all  chlorides  react  with  sulphuric  acid  and  are  converted 
into  sulphates,  the  reaction  indicated  below  takes  place: 

MnCl2  +  H2SO4  =  2HC1  +  MnSO4 

The  three  equations  given  above  represent  what  occurs  when 
manganese  dioxide,  sodium  chloride,  and  sulphuric  acid  are  mixed. 
The  products  obtained  are  sodium  sulphate,  Na2SO4,  manganese 
sulphate,  MnSO4,  water,  and  chlorine.  Hydrochloric  acid  and 
manganese  chloride,  MnCl2,  are  so-called  intermediate  products; 
they  react  with  other  substances  as  soon  as  formed — a  fact  indi- 
cated by  the  appearance  of  their  formulas  on  the  left-hand  side  of 
the  equations;  they  are  not  final  products  and  their  formulas  do 
not  appear  in  the  final  equation. 

A  single  equation  can  be  written  to  express  the  sum  of  all 
these  reactions.  First  the  formulas  of  the  substances  used  are 
set  down,  then  an  arrow,  and  next  the  formulas  of  the  final  prod- 
ucts of  the  reaction,  thus : 

NaCl  +  H2SO4  +  MnO2  ->  Na2SO4  +  MnSO4  +  H2O  +  C12 

The  expression  is  next  balanced  beginning  with  the  element 
sodium,  Na.  This  is  done  by  putting  a  2  in  front  of  the  formula 
NaCl,  since  two  sodium  atoms  are  present  in  sodium  sulphate, 
Na2SO4.  Having  balanced  the  sodium  in  sodium  sulphate  we 
next  balance  the  group  of  atoms,  SO4,  which  is  present  in  this 
compound,  and  also  in  manganese  sulphate  and  in  sulphuric  acid; 
this  is  done  by  taking  two  molecules  of  sulphuric  acid ;  the  expres- 
sion becomes  as  the  result  of  these  changes, 

2NaCl  +  2H2SO4  +  MnO2  ->  Na2SO4  +  MnSO4  +  H2O  +  C12 


104  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Proceeding  with  the  acid  in  which  the  SO4  group  has  been  bal- 
anced, we  next  examine  the  number  of  hydrogen  atoms;  these  are 
balanced  by  putting  2  in  front  of  the  formula  of  water.  The 
other  atom  present  in  this  compound,  oxygen,  is  next  balanced; 
there  are  now  two  atoms  on  the  right  of  the  arrow  and  two  on  the 
left  in  combination  with  manganese.  As  no  change  is  necessary 
at  this  step  we  balance  the  manganese;  no  change  is  required 
here.  The  remaining  atoms,  chlorine,  are  examined  and  found 
also  to  be  balanced.  The  expression  becomes  as  the  result  of  all 
these  changes, 

2NaCl  +  2H2SO4  +  MnO2  =  Na2SO4  +  MnSO4  +  2H2O  +  C12 

The  arrow  has  been  replaced  by  the  sign  of  equality,  since  a 
re-examination  of  the  entire  equation  shows  that  all  atoms  are 
balanced. 

The  student  should  carefully  study  the  process  by  which  the 
equation  was  balanced.  It.  is  of  the  greatest  importance  that 
facility  should  be  gained  in  carrying  out  the  process,  for  when 
this  is  acquired,  the  writing  of  chemical  equations  offers  little 
difficulty. 

115.  There  are  a  number  of  substances  besides  manganese 
dioxide  which  react  with  hydrochloric  acid  and  produce  chlorine. 
One  of  these,  potassium  permanganate,  can  be  used  conveniently 
for  this  purpose;   if  concentrated  hydrochloric  acid  is  allowed  to 
fall  on  it,  drop  by  drop,  at  room  temperature,  the  gas  is  formed. 
The  cost  of  potassium  permanganate  interferes  with  the  general 
use  of  this  way  of  making  chlorine. 

116.  (d)  Bleaching  Powder  and  Acids. — Chlorine  can  be  pre- 
pared conveniently  by  treating  bleaching  powder  with  an  acid. 
Bleaching  powder  is  made  by  treating  slaked  lime  with  chlorine. 
The  formula  of  lime,  calcium  oxide,  is  CaO,  and  that  of  bleaching 
powder  CaOCl2;  for  this  reason  it  is  often  called  chloride  of  lime; 
it  has  important  uses  which  will  be  described  later.     The  reactions 
which  take  place  between  this  substance  and  hydrochloric  acid 
and  sulphuric  acid,  respectively,  are  represented  by  the  following 
equations : 

CaOCl2  +  2HC1     =  CaCl2   +  H2O  +  C12 
CaOCl2  +  H2S04  =  CaS04  +  H2O  +  C12 


CHLORINE.     VALENCE  105 

In  the  first  case  calcium  chloride  is  formed,  along  with  water  and 
chlorine:  in  the  second,  calcium  sulphate.  It  is  to  be  noted  that 
all  substances  which  contain  the  group  of  atoms  represented  by 
SO4  are  called  sulphates;  applying  this  statement  to  sulphuric 
acid,  H2SO4,  it  should  be  called  hydrogen  sulphate — a  name  which 
is  frequently  used  for  the  acid. 

Chloride  of  lime  is  readily  available,  as  it  can  be  purchased  at 
grocery  stores.  For  this  reason  it  is  a  convenient  source  of  chlo- 
rine if  it  is  desired  to  use  it  in  the  household  for  bleaching  or 
other  purposes.  Vinegar  contains  an  acid,  acetic  acid,  and  can 
be  used  to  liberate  the  chlorine. 

117.  Properties  of  Chlorine. — Chlorine  is  a  heavy,  greenish- 
yellow  gas  which  has  a  stifling  odor.     It  affects  seriously  the 
lining  of  the  throat,  nose,  and  lungs,  and  when  inhaled  even  in 
moderate  quantities  produces  the  effects  which  result  from  a  bad 
cold.     In  larger  amounts  it  produces  suffocation  and  death.     On 
account  of  the3e  effects  chlorine  was  used  in  the  recent  war  to 
disable  an  army  before  an  attack  was  made.     The  gas  can  be 
liquefied, and  when  stored  in  cylinders  of  iron  is  readily  transported. 
When  the  valve  of  a  cylinder  containing  liquid  chlorine  is  opened 
the  substance  rushes  out  in  the  gaseous  condition.     As  it  is  a  very 
heavy  gas  it  stays  near  the  surface  of  the  earth  and  when  impelled 
by  the  wind  it  rolls  on  as  a  deep,  yellowish-green  cloud.     Certain 
compounds  prepared  from  chlorine  and  other  substances  affect  the 
eyes,  causing  intense  pain  and  a  copious  flow  of  tears;   such  sub- 
stances were  used  in  warfare  to  help  rout  the  enemy.     In  order 
to  withstand  gas-attacks  all  soldiers  were  furnished  with  rubber 
gas-masks,  which  fitted  the  face  closely  and  were  connected  with 
a  cannister  that  contained  substances  to  absorb  the  gases.     Espec- 
ially prepared  charcoal  which  absorbs  into  its  pores  large  quanti- 
ties of  gas  was  used  for  this  purpose,  together  with  a  mixture  con- 
taining lime  and  sodium  hydroxide  which  reacts  with  chlorine  and 
the  other  gases  containing  this  element. 

118.  One  liter  of  chlorine  at  0°  and  760  mm.  pressure  weighs 
3.220  grams.     It  is  2.49  times  as  heavy  as  air.     Liquid  chlorine 
has  the  specific  gravity  1.41  at  20°;    it  boils  at   —  33.6°  and 
changes  to  a  solid  at  —  102°.     The  pressure  of  liquid  chlorine 
when  stored  in  cylinders  is  6.6  atmospheres  at  20°.     When  chlo- 
rine is  passed  into  water  at  20°,  215  volumes  of  the  gas  dissolve 


106  INORGANIC  CHEMISTRY  FOR  COLLEGES 

in  100  volumes  of  the  liquid;  the  solubility  at  0°  is  461  volumes 
in  100;  for  this  reason  the  gas  is  not  collected  in  the  laboratory 
over  water,  in  the  way  used  in  the  case  of  hydrogen  and  oxygen. 
It  is  collected  by  the  upward  displacement  of  air;  when  chlorine 
is  conducted  into  a  bottle  filled  with  air,  the  heavy  gas  settles  to 
the  bottom  and  as  it  rises  forces  the  air  out. 

119.  Chemical  Conduct  of  Chlorine.  With  Phosphorus.— The 
chemical  conduct  of  chlorine  can  be  best  illustrated  by  a  number 
of  striking  experiments.  Several  jars  of  the  gas  are  provided. 
Into  one  is  introduced  a  bit  of  phosphorus;  it  soon  reacts  with  the 
production  of  light.  On  account  of  this  fact  we  say  that  phos- 
phorus burns  in  chlorine.  The  definition  of  the  word  burn  has 
been  broadened  to  include  all  chemical  phenomena  which  take 
place  with  the  evolution  of  light  and  heat;  it  is  not  now  restricted 
to  a  change  involving  oxygen.  When  phosphorus  burns  in  an 
atmosphere  of  chlorine,  phosphorus  pentachloride,  a  yellow  solid, 
is  formed : 

2P  +  5C12  =  2PC1* 

It  will  be  recalled  that  it  yielded  phosphorus  pentoxide  when 
burned  in  oxygen. 

With  Antimony. — If  powdered  antimony  is  dropped  into  a  jar 
containing  chlorine  the  two  elements  react  and  a  brilliant  flash  is 
seen  where  each  particle  of  the  powder  unites  with  the  chlorine. 
Antimony  trichloride  is  formed: 

2Sb  +  3C12  =  2SbCl3 

When  antimony  is  burned  in  oxygen,  antimony  trioxide,  Sb2O3,  is 
the  result  of  the  reaction. 

With  Copper. — When  a  piece  of  copper  in  the  form  of  a  thin 
foil  is  gently  heated  and  then  introduced  into  chlorine,  union 
takes  place  with  the  evolution  of  light,  and  copper  chloride  is 
formed : 

Cu  +  C12  =  CuCl2 

When  copper  is  heated  in  oxygen,  copper  oxide,  CuO,  is  produced. 
With  Sodium. — Sodium,  as  we  have  seen,  is  a  very  active 
metal  which  decomposes  water  with  great  violence;  chlorine  is  a 
poisonous  gas.  When  these  substances  unite  we  get  as  the  result 
sodium  chloride,  common  table-salt.  The  experiment  is  an  easy 


CHLORINE.     VALENCE  107 

one  to  carry  out.  Some  of  the  metal  is  rolled  into  a  thin  sheet, 
warmed  a  little,  and  put  into  a  jar  of  chlorine;  it  burns  and  a  white 
powder  is  formed  which  is  recognized  by  its  taste.  The  equation 
for  the  reaction  is, 

2Na  +  C12  =  2NaCl 

120.  With  Hydrogen. — When  a  jet  of  burning  hydrogen  is  put 
into  a  jar  of  chlorine  it  continues  to  burn,  this  time,  however,  with 
a  luminous  flame,  until  the  chlorine  is  used  up — a  fact  that  can 
be  recognized  by  the  disappearance  of  the  yellow  color.     The 
product  is  a  colorless  gas,  hydrogen  chloride,  which,  when  dissolved 
in  water,  is  called  hydrochloric  acid: 

H2  +  C12  =  2HC1 

The  fact  that  an  acid  is  formed  can  be  shown  by  testing  the  gas 
with  a  moist  piece  of  blue  litmus  paper,  which  changes  to  pink  in 
the  presence  of  acids. 

Hydrogen  and  chlorine  can  be  made  to  unite  through  the 
action  of  light.  If  a  thin- walled  glass  vessel  is  filled  with  a  mix- 
ture of  these  gases  and  placed  in  the  direct  sunlight,  the  union 
takes  place  with  explosive  violence.  Chlorine  and  many  of  its 
compounds  are  susceptible  'to  the  action  of  light;  we  shall  see 
later  that  a  number  of  photographic  processes  are  based  upon 
this  fact.  When  a  flame  is  introduced  into  a  mixture  of  hydrogen 
and  chlorine  the  gases  unite  with  the  production  of  a  loud  noise. 
The  behavior  of  hydrogen  with  chlorine  recalls  in  many  ways  the 
action  of  this  element  with  oxygen. 

121.  With  Other  Elements. — Chlorine  reacts  with  nearly  all 
the  elements,  chlorides  being  formed  more  or  less  rapidly  at  the 
ordinary  temperatures.     Carbon  resists  the  action  of  chlorine  and 
is  employed  in  making  electrodes  to  be  used  in  the  preparation  of 
the  gas  by  electrolytic  methods.     If  chlorine  and  the  element  with 
which  it  is  to  react  are  both  carefully  dried  before  being  brought 
together,  the  action  takes  place  very  slowly  in  many  cases.     Thus, 
bright  metallic  sodium  can  be  kept  without  change  in  an  atmos- 
phere of  dry  chlorine  for  a  long  time.     Water  in  this  case  is  a 
catalyst,  which  markedly  affects  the  rate  at  which  the  chemical 
change  takes  place.     Water  often  acts  in  this  way,  for  many  sub- 
stances which  react  rapidly,-  and  sometimes  with  violence,  under 


108  INORGANIC  CHEMISTRY  FOR  COLLEGES 

ordinary  conditions,  do  not  appear  to  affect  each  other  if  water  is 
rigorously  excluded. 

122.  With  Hydrocarbons. — The  experiments  with  hydrogen 
and  chlorine  which  have  been  described  show  that  these  elements 
possess  a  great  affinity  for  each  other — the  chemical  energy  bound 
up  with  each  element  is  of  such  a  nature  that  it  tends  to  escape  as 
heat  when  the  opportunity  is  offered.  Not  only  does  chlorine 
react  with  free  hydrogen,  but  it  withdraws  hydrogen  from  com- 
pounds containing  it.  Hydrocarbons  belong  to  a  class  of  sub- 
stances which  are  composed  of  hydrogen  and  carbon — a  fact 
which  gives  them  their  name.  Turpentine  is  a  familiar  hydro- 
carbon which  is  obtained  from  pine  trees.  If  some  of  this  sub- 
stance is  warmed  and  poured  on  a  piece  of  paper  which  is  then 
put  into  a  jar  containing  chlorine,  a  striking  phenomenon  takes 
place;  a  dense  black  cloud  of  very  finely  divided  carbon  is  formed; 
and  a  test  of  the  contents  of  the  jar  will  show  that  hydrogen 
chloride  is  present.  The  reaction  consists  in  the  union  of  the  chlo- 
rine and  hydrogen,  and  the  carbon  is  left  behind.  Other  hydro- 
carbons and  compounds  containing  carbon,  hydrogen,  and  oxygen, 
react  in  a  similar  way.  When  a  burning  candle  is  lowered  into  a 
jar  of  chlorine,  carbon  is  deposited  as  it  burns,  and  hydrogen 
chloride  is  formed.  Other  compounds  are  produced,  however, 
and  the  phenomenon  is  not  so  striking  as  when  turpentine  is  used. 
When  most  hydrocarbons  and  compounds  which  contain  carbon, 
hydrogen,  and  oxygen  are  subjected  to  the  action  of  chlorine, 
hydrogen  chloride  is  formed,  and  chlorine  enters  the  molecule  and 
takes  the  place  of  the  displaced  hydrogen.  For  example,  when 
methane,  CH4,  which  is  the  chief  constituent  of  natural  gas,  is 
treated  with  chlorine  such  a  reaction  takes  place: 

CH4  +  4C12  =  CC14  +  4HC1 

This  particular  kind  of  reaction  is  called  substitution,  since  one 
element  takes  the  place  of  another.  Reactions  of  this  type  are 
greatly  facilitated  by  sunlight.  The  discovery  that  one  element 
could  take  the  place  of  another  in  this  way  was  the  result  of  an 
unusual  occurrence.  A  state  ball  was  being  held  at  the  Tuilleries 
in  the  reign  of  Louis  Philippe.  Soon  after  the  candles  were  lighted 
the  room  was  filled  with  suffocating  fames.  The  guests  were  dis- 
missed, and  the  cause  of  the  remarkable  occurrence  was  investi- 


CHLORINE.     VALENCE  109 

gated.  The  master  of  ceremonies  happened  to  be  a  relative  of  a 
young  chemist,  Dumas  by  name,  who  was  asked  to  find  out  the 
source  of  the  fumes.  He  found  that  the  candles  had  been  bleached 
by  chlorine,  and  that  some  of  the  element  had  driven  out  a  part  of 
the  hydrogen  in  the  wax  of  which  the  candles  were  made,  and  had 
taken  its  place.  When  the  candles  were  afterwards  burned  this 
chlorine  was  given  off  in  combination  with  hydrogen,  as  hydrogen 
chloride;  and  this  was  the  substance  which  caused  the  discom- 
fiture of  the  king  and  his  guests.  This  type  of  reaction  was  studied 
carefully  by  Dumas,  and  became  of  the  greatest  importance  in 
organic  chemistry. 

123.  With  Water. — If  water  through  which  chlorine  is  passing 
is  cooled  to  0°,  a  white,  crystalline  substance  separates.     It  is 
called  chlorine  hydrate,   and  has  the  formula  Ck,  8H20;    the 
comma  placed  between  the  formula  of  chlorine  and  that  of  water 
is  used  to  indicate  the  fact  that  the  compound  readily  decomposes 
into  the  constituents  of  which  it  is  made  up.     Chlorine  hydrate 
belongs  to  the   class  of  substances  often  called  molecular   com- 
pounds, i.e.,  compounds  formed  as  the  direct  union  of  two  or 
more  molecules,  which  more  or  less  readily  break  down  into  these 
same  molecules.     Water  forms  many  such  molecular  compounds, 
which    are    commonly    called    hydrates.     Sulphuric    acid,    for 
example,  forms  a  hydrate  for  which  the  formula  is  EkSCU,  H^O. 

124.  Michael  Faraday,  who  worked  in  the  Royal  Institution 
in  London,  a  place  where  many  great  discoveries  in  physics  and 
chemistry  have  been  made,  was  studying  (1823)  chlorine  hydrate 
when  he  hit  upon  a  method  of  liquefying  chlorine.     He  heated 
some  of  the  hydrate  in  a  closed  tube  shaped  like  an  inverted 
letter  V;  one  arm  of  the  tube  contained  the  hydrate  and  the  other 
was  surrounded  by  water.     A  yellow  oil  which  appeared  in  the 
cold  part  of  the  tube  proved  to  be  pure  chlorine  in  the  form  of  a 
liquid.     It  was  produced  as  the  result  of  the  fact  that  when  the 
chlorine  hydrate  was  heated  it  dissociated  into  its  constituents, 
and  the  gas  liberated  in  such  a  small  volume  created  a  great  pres- 
sure and,  as  a  consequence,  turned  to  a  liquid.     This  observation 
led  Faraday  to  study  other  gases,  and  he  finally  succeeded  in 
liquefying  a  number  by  subjecting  them  to  pressure  at  low  temp- 
eratures.    Many  years  after,  Dewar  took  up  the  work  again  in 
the  Royal  Institution  and  with  the  aid  of  improved  methods  and 


110  INORGANIC  CHEMISTRY  FOR  COLLEGES 

the  knowledge  gained  as  chemistry  and  physics  advanced,  liquefied 
the  gases  which  did  not  yield  to  Faraday's  attempts. 

125.  A  solution  of  chlorine  in  water  contains,  in  addition  to 
chlorine,  some  hydrochloric  acid  and  a  compound  of  the  com- 
position represented  by  the  formula  HOC1,  called  hypochlorous 
acid.  The  substances  are  formed  as  the  result  of  the  reaction 
represented  as  follows  : 

H2O  +  C12  =  HOC1  -4-  HC1 

In  this  case  one-half  the  hydrogen  in  the  molecule  of  water  is 
substituted  by  chlorine,  HOH  —  >  HOC1.  All  the  chlorine  does 
not  disappear  from  the  solution,  however,  because  hypochlorous 
acid  and  hydrochloric  acid  interact  to  form  water  and  chlorine, 
the  equation  for  the  reaction  being  the  reverse  of  that  just  given: 

HOC1  +  HC1  =  H2O  +  C12 

We  have  in  this  case  what  is  called  a  reversible  reaction  —  one  in 
which  certain  substances  interact  to  form  others  which,  in  turn, 
interact  to  form  the  original  substances.  This  is  expressed  briefly 
by  replacing  the  sign  of  equality  in  the  equation  by  the  symbol  ^±, 
thus: 


If  water  and  chlorine  are  brought  together  they  interact  to  form 
hypochlorous  acid  and  hydrochloric  acid,  and  these  in  turn  to 
form  chlorine  and  water;  the  solution  contains,  as  a  consequence, 
all  four  substances.  When  one  volume  of  chlorine  is  dissolved  in 
one  volume  of  water  at  10°  about  66  per  cent  of  the  chlorine  is 
present  as  free  chlorine  and  the  rest  as  hypochlorous  acid  and 
hydrochloric  acid.  This  fact  can  be  represented  thus: 

34 

H2O  +  C12  <=±  HOC1  +  HC1 
66 

If  a  solution  of  chlorine  in  water  (chlorine-water)  is  exposed  to 
direct  sunlight  the  hypochlorous  acid  present  decomposes  according 
to  the  following  equation: 

2HOC1  =  2HC1  -f  02 


CHLORINE.    VALENCE  111 

As  the  oxygen  is  slowly  set  free  it  separates  from  the  solution  in 
bubbles  and  can  be  collected  in  a  suitable  apparatus.  As  soon  as 
the  hypochlorous  acid  begins  to  decompose  more  chlorine  reacts 
with  the  water  present  to  form  the  acid;  this  process  continues 
until  all  the  chlorine  disappears,  provided  the  reaction  takes  place 
in  sunlight;  the  complete  change  which  takes  place  is  represented 
the  following  equation: 

2H2O  +  2C12  =  4HC1  +  02 

This  reaction  is  the  reverse  of  that  written  for  the  Deacon  process 
for  the  manufacture  of  chlorine;  it  is,  thus,  a  reversible  reaction, 
and  it  is  for  this  reason  that  when  oxygen  and  hydrogen  chloride 
react  they  do  not  change  completely  to  chlorine  and  water — a 
fact  which  has  been  noted.  The  reaction  which  takes  place  in  the 
Deacon  process  is  brought  about  when  the  substances  are  in  the 
gaseous  condition;  the  reaction  in  the  reverse  direction  proceeds 
best  when  chlorine  is  dissolved  in  water.  The  conditions  under 
which  a  reversible  reaction  takes  place  influence  greatly  the  extent 
to  which  it  proceeds  in  either  direction.  This  important  fact 
will  be  discussed  later. 

126.  Chlorine  as  an  Oxidizing  Agent. — When  substances  are 
treated  with  chlorine  and  water  they  are  oxidized,  the  oxygen 
for  the  purpose  being  furnished  by  the  hypochlorous  acid  present 
in  the  solution.  We  have  not  as  yet  become  familiar  with  such 
substances,  but  a  simple  case  can  be  cited  as  an  illustration. 
Sulphurous  acid  has  the  formula  H^SOs;  it  can  be  changed  by 
oxygen — oxidized — to  sulphuric  acid,  H2SO4.  The  changes  which 
occur  when  sulphurous  acid  is  treated  with  chlorine  and  water  can 
be  expressed  by  the  following  equations: 

C12  +  H2O  =  HOC1  +  HC1 

HOC1  =  HC1  +  O 
H2SO3  +  O=H2SO4 

These  equations  can  be  combined  into  one  by  the  process  which  has 
been  described  at  length  (71) : 

C12  +  H2O  +  H2SO3  =  2HC1  +  H2S04 


112  INORGANIC  CHEMISTRY  FOR  COLLEGES 

127.  It  should  be  noted  that  we  do  not  write  the  second 
equation  above  in  the  form  used  before : 

2HOC1  =  2HC1  +  O2 

The  formula  for  oxygen  gas  is  02.  When  chlorine  and  water  act  as 
an  oxidizing  agent  no  oxygen  gas  is  given  off;  we  assume,  therefore, 
that  as  soon  as  an  oxygen  atom  is  set  free  it  immediately  unites 
with  the  sulphurous  acid  and  forms  sulphuric  acid  according  to 
the  third  equation  above.  There  is  experimental  evidence  in 
favor  of  this  view.  When  oxygen  is  formed  in  the  presence  of  a 
substance  that  can  be  oxidized,  it  is  more  reactive  than  oxygen 
gas.  Many  substances  which  do  not  react  with  the  latter  at 
ordinary  temperatures,  do  react  with  oxygen  if  the  element  is 
liberated  from  a  compound  in  the  presence  of  these  substances. 
In  order  to  make  this  distinction,  we  say  that  the  element  when 
liberated  from  a  compound  is  in  the  nascent  state,  the  word  nas- 
cent being  derived  from  the  Latin  word  nascens,  born.  Chlorine 
and  water  furnish,  thus,  a  convenient  source  of  nascent  oxygen. 
The  explanation  of  these  facts  from  the  theory  of  atoms  is  that  the 
oxygen  when  set  free  is  in  the  form  of  atoms — the  symbol  for  nas- 
cent oxygen  is  O ;  if  nothing  is  present  to  combine  with  these  atoms 
they  unite  and  form  molecules,  and  thus  produce  oxygen  gas, 
the  formula  for  which  is  62 .  The  activity  of  nascent  oxygen  is 
explained  also  from  the  standpoint  of  energy.  When  the  atoms 
are  set  free  they  possess  a  large  amount  of  chemical  energy  and 
are,  therefore,  active.  If  nothing  is  present  to  be  oxidized  these 
atoms  lose  a  part  of  their  energy  when  they  react  with  each  other 
to  form  molecules.  This  reaction  can  be  expressed  thus: 

2O  =  O2 

When  it  takes  place  a  large  amount  of  chemical  energy  is  lost  and 
the  resulting  molecule  is.  accordingly,  less  active. 

128.  Similar  relationships  are  observed  in  the  case  of  hydrogen. 
Many  substances  that  are  not  affected  by  hydrogen  gas  are  re- 
duced when  placed  in  a  vessel  in  which  hydrogen  is  being  formed 
as  the  result  of  the  action  of  an  acid  on  a  metal.     A  similar  explana- 
tion is  offered.     Atoms  of  hydrogen  are  first  liberated;   these  are 


CHLORINE.    VALENCE  113 

active;  if  no  reducible  substance  is  present  they  combine,  lose  a 
part  of  their  chemical  energy,  and  the  resulting  molecule  is  less 
active  than  the  atoms  of  which  it  is  composed.  What  occurs 
according  to  this  view  when  zinc  and  hydrochloric  acid  react  can 
be  expressed  by  equations  as  follows: 

Zn  +  2HC1  =  ZnCl2  +  2H 
2H  =  H2 

The  hypothesis  put  forward  to  account  for  the  nascent  state  is 
valuable  since  it  offers  an  explanation  of  facts  other  than  those 
which  led  to  its  suggestion.  It  will  be  recalled  that  gaseous  hydro- 
gen does  not  reduce  copper  oxide  at  room  temperature  to  copper 
and  water;  the  substances  must  be  heated  together.  It  is  highly 
probable  that  at  the  higher  temperature  the  molecules  of  hydrogen 
break  down  in  part  into  atoms;  these,  according  to  the  hypothesis, 
are  more  active  than  molecules,  and,  as  a  consequence,  they  with- 
draw the  oxygen  from  the  copper  oxide.  If  this  view  is  correct 
the  molecules  of  hydrogen  first  dissociate  into  atoms;  the  reaction 
which  takes  place  can  be  represented  as  follows; 

H2  =  2H 

It  has  been  shown  that  many  molecules  can  be  decomposed  by  heat 
into  parts  which  reunite  when  the  source  of  heat  is  removed; 
when  this  decomposition  takes  place  the  molecules  are  said  to 
dissociate.  Water,  for  example,  dissociates  to  the  extent  of  1.8 
per  cent  at  2000°  into  hydrogen  and  oxygen.  If  the  mixture  of 
the  three  substances  at  this  temperature  is  cooled,  the  hydrogen 
and  oxygen  recombine  to  form  water. 

129.  Chlorine  as  a  Bleaching  Agent. — If  pieces  of  dyed  cotton 
cloth  which  have  been  previously  dried  are  suspended  in  a  jar  of 
dry  chlorine,  scarcely  any  change  takes  place.  If  other  pieces  of 
the  same  cloth  are  dipped  in  water  and  then  subjected  to  the  action 
of  chlorine  in  the  same  way,  a  change  is  soon  observable;  the  dyes 
begin  to  fade,  and  after  a  few  minutes  the  pieces  of  cloth  become 
white  or  nearly  so.  These  experiments  show  that  chlorine  and 
water  destroy  certain  dyes.  They  affect  in  a  similar  way  the 
coloring  matter  present  in  unbleached  cotton,  and  could  be  used 


114  INORGANIC  CHEMISTRY  FOR  COLLEGES 

for  bleaching.  The  reaction  which  takes  place  is  expressed,  in  the 
main,  by  the  equations  already  given: 

H2O  +  C12  =  HOC1  +  HC1 
HOC1  =  HC1  +  O 

The  nascent  oxygen  then  oxidizes  the  colored  compound,  which,  as 
a  result,  is  converted  into  other  substances  that  are  colorless. 

130.  Test  for  Chlorine. — When  free  chlorine  is  brought  into 
contact  with  a  solution  of  potassium  iodide,  KI,  the  iodine  is  set 
free  and  potassium  chloride  is  formed: 

2KI  +  C12  =  2KC1  +  I2 

The  iodine  can  be  recognized  even  when  it  is  present  in  very  small 
amounts  by  shaking  the  solution  with  carbon  disulphide;  it  dis- 
solves in  the  latter  and  imparts  to  it  a  characteristic  purple 
color.  The  iodine  can  also  be  detected  by  treating  it  with  a  solu- 
tion made  by  heating  a  little  starch  with  water;  in  this  case 
a  blue  color  is  produced.  These  reactions  can  be  used  to  detect 
chlorine,  but  since  other  substances  liberate  iodine  from  potassium 
iodide  they  alone  do  not  prove  the  presence  of  the  free  element. 
In  order  to  draw  a  definite  conclusion  the  solution  is  tested  for  the 
chloride  formed,  by  the  method  described  in  section  145. 

131.  Uses  of  Chlorine. — Most  of  the  chlorine  that  is  manu- 
factured is  converted  into  chloride  of  lime,  CaOCk,  or  sodium 
hypochlorite,  NaOCl,  compounds  which  are  extensively  used  in 
bleaching.     Chloride  of  lime  yields  chlorine  readily  and  is  used  as  a 
disinfectant,  a  deodorizer,  in  the  manufacture  of  chloroform,  and 
for  other  purposes.     Free  chlorine  in  water  destroys  bacteria; 
the  gas  is,  accordingly,  now  used  in  sterilizing  water  supplies  when 
contamination  is  suspected  or  has  been  shown  to  be  present. 
Chlorine  is  used  for  making  chlorides  of  many  of  the  elements,  in 
the  preparation  of  potassium  chlorate,  and  in  separating  other 
elements  from  their  compounds,  such  as  bromine  and  iodine. 

Tin  is  recovered  from  waste  tin  cans  by  an  ingenious  method 
which  involves  the  use  of  chlorine.  Tin  cans  are  made  from 
sheet-iron  which  has  been  covered  with  tin.  As  the  latter  is  a 
more  or  less  expensive  metal  it  is  desirable  to  recover  it  from  used 
cans.  The  problem  was  finally  solved  in  a  simple  way.  When 


CHLORINE.     VALENCE  115 

chlorine  was  passed  over  the  cans,  the  tin  was  converted  into  the 
chloride,  SnCU,  which  boils  at  114°.  The  mixture  was  then 
heated  and  the  chloride  distilled  off.  The  product,  stannic  chloride, 
is  used  in  large  quantities  as  a  mordant  in  dyeing  silk. 

Chlorine  is  also  used  in  the  preparation  of  carbon  tetrachloride, 
a  substance  which  has  recently  been  shown  to  be  very  useful 
on  account  of  the  fact  that  it  is  a  liquid  that  readily  passes  into 
a  vapor  which  is  non-inflammable  and,  the  presence  of  which  pre- 
vents combustible  substances  from  burning.  It  is  used  in  fire 
extinguishers,  and  when  added  to  gasoline  makes  a  mixture  that 
can  be  used  for  cleaning  purposes  without  the  danger  which  attends 
the  use  of  gasoline  alone. 

132.  Comparison  of  the  Chemical  Conduct  of  Chlorine  with 
that  of  Oxygen. — The  acquisition  of  the  large  number  of  facts  that 
one  meets  in  the  study  of  chemistry  is  greatly  facilitated  by  exam- 
ining new  facts  as  they  are  presented  in  the  light  of  the  knowledge 
already  gained.  This  is  best  done  by  searching  out  resemblances 
and  differences,  and  by  determining  whether  the  new  facts  are 
additional  examples  of  general  principles  that  have  been  learned. 
This  method  will  be  illustrated  by  an  examination  of  the  chem- 
istry of  oxygen  and  chlorine.  First  resemblances  will  be  noted, 
then  differences. 

Occurrence:  Both  elements  are  active  and  occur  in  combina- 
tion with  other  elements;  oxides  and  chlorides  of  metals  are 
important  substances  found  in  nature.  A  much  larger  proportion 
of  oxygen  than  of  chlorine  is  present  in  the  earth;  oxygen  is  found 
in  the  free  condition,  chlorine  is  not. 

Preparation:  Both  elements  can  be  obtained  by  heating  the 
compounds  which  they  form  with  inactive  metals;  platinum  oxide 
yields  oxygen  and  platinum  chloride  yields  chlorine  when  heated. 
If  the  metal  in  combination  with  chlorine  or  oxygen  is  an  active 
one,  such  a  decomposition  does  not  take  place;  we  cannot  decom- 
pose sodium  chloride  or  sodium  oxide  by  heat.  Some  chlorides 
decompose  more  readily  than  the  corresponding  oxides,  and  some 
do  not;  manganese  tetrachloride  breaks  down  spontaneously  into 
manganese  dichloride  and  chlorine;  in  order  to  decompose  man- 
ganese dioxide  a  high  temperature  is  required.  On  the  other  hand, 
mercuric  oxide  is  decomposed  into  mercury  and  oxygen  at  a 
lower  temperature  than  that  at  which  an  analogous  decomposi- 


116  INORGANIC  CHEMISTRY  FOR  COLLEGES 

tion  takes  place  in  the  case  of  mercuric  chloride.  Both  elements 
can  be  prepared  by  the  action  of  an  electric  current  on  their  com- 
pounds with  hydrogen;  hydrogen  oxide,  water,  under  these  cir- 
cumstances yields  hydrogen  and  oxygen,  and  hydrogen  chloride 
yields  hydrogen  and  chlorine. 

Chemical  conduct :  The  two  elements  unite  with  most  other  ele- 
ments and  form  oxides  and  chlorides.  The  phenomena  observed  in 
the  two  cases  resemble  each  other  very  closely;  hydrogen  burns  in 
oxygen,  also  in  chlorine;  phosphorus  and  other  elements  behave 
in  a  similar  way;  but  carbon  burns  in  oxygen  and  not  in  chlorine. 
The  kindling  temperature  for  a  reaction  between  an  element  and 
chlorine  is  usually  lower  than  that  of  the  analogous  reaction 
with  oxygen.  Most  substances  react  with  chlorine  at  room  tem- 
perature more  rapidly  than  with  oxygen.  In  general,  chlorine  is 
the  more  active  element  at  ordinary  temperatures.  Chlorine 
withdraws  hydrogen  from  compounds  containing  it,  and  unites 
with  the  hydrogen  and  also  with  the  element  to  which  the  hydro- 
gen was  joined.  Oxygen  behaves  in  a  similar  way.  Equations 
for  two  typical  reactions  are  as  follows : 

CH4  +  4C12  =  CC14  +  4HC1 
CH4  +  2O2  =  CO2  +  2H2O 

It  will  be  seen  from  the  comparisons  which  have  been  drawn 
that  chlorine  and  oxygen  resemble  each  other  very  closely  in 
chemical  properties.  The  facts  which  have  been  given  in  the 
discussion  of  the  two  elements  can  be  remembered  more  readily 
and  appreciated  more  fully  as  the  result  of  this  comparison.  At 
this  stage  in  his  study  of  chemistry  the  student  would  probably 
not  be  able  to  make  as  full  a  comparison  as  that  given  above,  but 
as  he  advances  he  will  soon  learn  to  generalize  his  facts  if  he  con- 
stantly examines  carefully  new  facts  as  they  are  presented.  He 
should  get  in  the  habit  of  asking,  is  this  like  anything  I  have 
learned  before,  can  I  hang  this  new  fact  on  an  old  one?  Of  course 
this  most  efficient  method  of  study  can  be  used  only  when  the 
subject  is  studied  from  day  to  day;  if  the  facts  are  not  acquired 
as  they  are  presented,  they  cannot  evidently  be  used  to  help 
remember  new  facts.  A  real,  usable  knowledge  of  chemistry  can 
be  gained  only  through  continuous  study;  in  no  other  way  can 
one  get  the  benefit  that  is  won  through  the  study  of  a  science. 


CHLORINE.     VALENCE  117 

133.  Valence. — A  comparison  of  the  formulas  of  a  series  of 
compounds  containing  chlorine  with  those  of  a  series  of  com- 
pounds of  the  same  elements  with  oxygen,  brings  out  some  striking 
relations.  Let  us  examine  a  few  of  these  set  side  by  side: 


HC1 

H2O 

NaCl 

Na20 

ZnCl2 

ZnO 

CaCl2 

CaO 

A1C13 

A1203 

SbCl3 

Sb2O3 

MnCU 

Mn02 

ecu 

CO2 

We  see  from  the  first  column  that  1  hydrogen  atom  unites  with 
1  chlorine  atom,  and  that  1  sodium  atom  unites  with  1  chlorine 
atom;  zinc  and  calcium  each  unite  with  2  chlorine  atoms;  alumi- 
nium and  antimony  with  3;  and  manganese  and  carbon  with  4. 
The  elements  differ  in  the  number  of  atoms  of  chlorine  which  they 
can  hold  in  combination.  An  examination  of  the  second  column 
shows  not  only  that  these  same  elements  differ  in  their  power  to 
hold  oxygen  atoms,  but  that  there  is  remarkable  relationship  be- 
tween the  combining  power  of  any  element  as  measured  by  chlor- 
ine in  one  case  and  by  oxygen  in  the  other.  This  will  be  clear 
from  the  following  considerations :  Hydrogen  is  taken  as  the  stand- 
ard of  combining  power  when  we  wish  to  consider  the  number  of 
atoms  which  unite  with  one  another.  We  say  the  combining 
power  of  hydrogen  is  1  or  use  a  more  technical  word  and  say  its 
valence  is  1,  the  word  being  derived  from  the  Latin  word  valentia 
signifying  strength.  Since  1  chlorine  atom  unites  with  1  hydrogen 
atom  its  valence  is  also  1.  We,  thus,  have  two  unit  standards  by 
which  we  can  measure  the  valence  of  any  other  element. 

We  can  now  readily  express  the  differences  in  combining  power 
observed  among  the  elements  which  occur  in  the  formulas  listed 
in  the  first  column  above.  Hydrogen  and  sodium  each  have  the 
valence  1,  zinc  and  calcium  have  the  valence  2,  aluminium  and 
antimony  have  the  valence  3,  and  manganese  and  carbon  the 
valence  4.  These  conclusions  are  drawn  from  the  study  of  the  for- 
mulas of  the  chlorides  of  these  elements.  We  shall  next  examine 
the  oxides.  One  oxygen  atom  unites  with  2  hydrogen  atoms;  it, 


118  INORGANIC  CHEMISTRY  FOR  COLLEGES 

accordingly,  has  the  valence  2.  Knowing  this  we  can  determine 
the  valence  of  the  elements  in  the  compounds  in  the  second  column 
by  comparing  them  with  oxygen  and  remembering  that  its  valence 
is  2.  In  the  formula  Na2O,  2  sodium  atoms  are  represented  in 
combination  with  1  oxygen  atom ;  the  combining  power  of  the  lat- 
ter which  is  2  is  used  up  in  holding  2  sodium  atoms;  each  of  the  lat- 
ter must  have  a  combining  power  of  1 — the  valence  of  sodium  is  1 . 

In  zinc  oxide,  ZnO,  the  valence  of  zinc  is  2  because  oxygen  has 
this  valence,  and  1  atom  is  in  union  with  1  atom.  It  must  not  be 
concluded  that  the  valence  of  zinc  is  1  because  it  unites  with  1 
oxygen  atom.  We  must  always  go  back  to  the  standard  selected, 
namely  hydrogen,  and  in  this  standard  the  combining  power  of 
oxygen  is  2,  as  has  been  stated;  therefore,  that  of  zinc  is,  also  2. 
Calcium  is  like  zinc  in  valence. 

When  we  come  to  aluminium  oxide  the  case  appears  to  be  more 
difficult.  It  is  simple  if  looked  at  in  the  following  way:  As  each 
oxygen  atom  has  the  combining  power  of  2,  3  atoms  of  this  element 
have  the  combining  power  of  6;  this  is  used  in  holding  in  combina- 
tion 2  atoms  of  aluminium;  and  each  atom  of  this  element  must 
have,  as  a  consequence,  the  valence  3;  likewise  antimony  has 
the  valence  3.  Manganese  and  carbon  each  unite  with  2  oxygen 
atoms  and,  consequently,  they  have  the  valence  4. 

As  a  result  of  the  study  of  the  formulas  of  the  chlorides  of  a 
series  of  elements,  we  come  to  a  conclusion  as  to  the  combining 
power  of  these  elements;  a  study  of  the  formulas  of  the  oxides 
leads  to  conclusions  as  to  their  combining  capacity  identical  with 
those  arrived  at  in  the  first  way.  This  is  a  striking  fact,  and  is 
more  remarkable  when  we  find  that  the  valence  which  each  of 
these  elements  possesses  in  other  compounds  is  the  same  as  that 
arrived  at  from  a  consideration  of  the  formulas  of  their  chlorides. 
The  valence  of  an  element  is  thus  an  expression  of  a  chemical 
characteristic  of  that  element.  We  shall  see  that  this  fact  sim- 
plifies markedly  remembering  chemical  formulas;  if  we  know  the 
valence  of  the  elements  which  are  present  in  a  compound  we  can, 
in  most  cases,  write  its  formula.  A  few  examples  will  make  this 
clear.  Zinc  always  has  the  valence  2.  What  will  be  the  for- 
mulas of  the  compounds  of  zinc  with  bromine,  sulphur,  and  phos- 
phorus? The  symbols  for  these  elements  and  the  valence  of  each 

i    ii  in 
are  represented  thus — Br,  S,  P.    A  Roman  numeral  is  often  placed 


CHLORINE.    VALENCE  119 

over  a  symbol  to  represent  the  valence  of  the  element.  When  the 
writing  of  chemical  formulas  becomes  "  second  nature  "  these 
numbers  can  be  omitted;  but  until  the  student  is  sure  of  himself, 
he  should  make  it  a  practice  to  indicate  the  valence  of  the  elements, 
for  such  a  practice  will  prevent  many  mistakes.  The  formulas 
of  the  compounds  of  the  elements  with  zinc  are  evidently  as  follows: 

ii    i  ii  ii  ii  in 

ZnBr2  ZnS  Zn3P2 

The  sum  of  the  combining  capacities  of  all  of  the  atoms  of  one 
element  in  the  compound  must  equal  the  sum  of  the  combining 
capacities  of  all  of  the  atoms  of  the  other.  In  the  case  of  zinc 
bromide,  1  zinc  atom  with  the  valence  2,  1  X  2  =  2,  is  united  with 
2  atoms  of  bromine  each  having  the  valence  1,  2X1=2; 
in  the  second  case,  zinc,  1X2=2,  and  sulphur,  1X2=2; 
in  the  third  compound  there  are  3  atoms  of  zinc,  3X2=6, 
and  2  atoms  of  phosphorus,  2X3=6.  As  a  mathematical 
check  on  the  correctness  of  a  formula  we  can  multiply  the  number 
of  atoms  of  one  element  present  by  its  valence,  and  the  number  of 
atoms  of  the  second  element  by  its  valence,  and  we  must  get  the 

same  number  in  each  case.     It  is  evident  that  the  following  do 

ii   i        ii   ii  ii  ii 

not  represent  chemical  compounds:  ZnCls,  Z^Ss  and  ZnaC^. 
Another  way  to  check  up  the  correctness  of  a  formula  is  to  repre- 
sent each  combining  power  by  a  line  drawn  from  the  symbol  of 
the  element;  the  number  of  lines  will,  evidently,  be  equal  to  the 
valence  of  the  element.  A  number  of  formulas  written  in  this 
way  will  be  self-explanatory: 

^ 

Nax  AlC 

>0  Zn  =  0  0 

Na/ 

Formulas  written  in  the  way  illustrated  above  are  called  graphic 
formulas;  they  are  useful  for  other  purposes  than  the  one  explained 
here.  In  representing  compounds  by  graphic  formulas  it  is  impor- 
tant to  emphasize  the  fact  that  there  can  be  no  free  combining 
powers — the  lines  must  run  from  the  symbol  of  one  element  to 
that  of  another.  Such  combinations  as  the  following  do  not  rep- 
resent compounds: 

Na,  / 

">O  AJ  =  0 


120  INORGANIC  CHEMISTRY  FOR  COLLEGES 

From  the  above  it  will  be  seen  that  much  mental  effort  will  be 
saved  if  the  valence  of  an  element  is  learned  as  soon  as  its  com- 
pounds are  met  with,  for  if  this  is  done  it  will  not  be  necessary  to 
memorize  a  large  number  of  formulas. 

134.  The    valence    of    certain    elements    is    constant;     thus, 
hydrogen,   sodium,  and  potassium   always  have  the  valence   1; 
zinc,    calcium,    and    magnesium,    2;    aluminium,    3,    etc.      The 
valence  of  other  elements  varies ;  iron,  for  example,  has  the  valence 
2  in  some  compounds  and  3  in  others ;  it  forms  two  chlorides  which 
have  the  formulas  FeCk  and  FeCla,  respectively.    This  is  a  source 
of  complexity,  but  one  that  is  readily  mastered  as  the  compounds 
are  studied. 

135.  The  method  of  writing  formulas  of  compounds  as  de- 
scribed above  is  applicable  in  the  case  of  those  which  contain  two 
elements  only.     It  can  be  applied  in  a  simple  way  to  other  com- 
pounds when  a  few  principles  are  made  clear.     We  have  seen  that 
the  formula  of  sulphuric  acid  is  B^SCU  and  that  of  sodium  sulphate 
is  Na2S(>4.     There  are  many  substances  known  which  can  be  pre- 
pared by  replacing  the  hydrogen  in  sulphuric  acid  by  atoms  of 
metals;    all  these  compounds  contain  the  group  of  atoms  repre- 
sented by  the  symbols  SO 4;    they  are  all  called  sulphates.     We 
have  calcium  sulphate  CaSCU,  copper  sulphate,  CuSO4,  zinc  sul- 
phate, ZnSC>4,  etc.     The  group  of  atoms  present  in  these  com- 
pounds, SO4,  is  called  a  radical,  the  word  being  derived  from  the 
Latin  word  radix,  meaning  root.     As  it  is  present  in  sulphuric  acid 
it  is  called  an  acid  radical.     We  can  assign  to  this  radical  a  valence; 
since  it  is  found  in  combination  with  2  hydrogen  atoms  in  sul- 
phuric acid,  H2SO4,  its  valence  is  2.     With  this  understood  it  is 
possible  to  write  the  formula  of  the  sulphate  of  any  metal  pro- 
vided the  valence  of  the  latter  is  known.     The  method  described 
in  detail  above  when  applied  to  writing  the  formulas  of  sodium 
sulphate,  zinc  sulphate,  and  aluminium  sulphate  leads  to  the  fol- 
lowing: 

i     ii  ii   ii  in     ii 

Na2S04       ZnSO4       A12(SO4)3 

In  this  case  the  valence  of  the  SO4  radical  is  written  above  it,  in 
the  way  used  with  the  elements.  In  the  formula  for  aluminium 
sulphate  the  radical  is  enclosed  in  a  parenthesis  and  a  subscript  3 
is  written  to  the  right  of  it;  this  indicates  that  3  radicals  each  with 


CHLORINE.     VALENCE  121 

the  valence  2,  3  X  2  =  6,  are  needed  to  combine  with  2  atoms 
of  aluminium  with  the  valence  3,  2  X  3  =  6.  In  order  to  write 
the  formulas  of  the  compounds  derived  from  acids  by  replacing 
the  hydrogen  of  the  latter  by  metals — compounds  which  are 
called  salts — it  is  necessary  to  know  the  valences  of  the  acid 
radicals. 

EXERCISES 

1.  Chlorine  can  be  prepared  by  the  action  of  hydrochloric  acid  on  lead 
dioxide,    PbO2,    and   on   potassium   permanganate,    KMnO4.     Convert   the 
following    into    balanced    chemical    equations:      (a)  PbO2  +  HC1  — >  PbCl2 
+  H2O  +  C12  and   (b)  KMnO4  +  HC1  -+  KC1  +  MnCl2  +  H2O  +  C12. 

2.  (a)  What  proportion  of  the  chlorine  in  sodium  chloride  is  set  free  when 
the  latter  is  treated  with  sulphuric  acid  and  manganese  dioxide?     (b)  What 
proportion  of  the  chlorine  is  set  free  when  hydrochloric  acid  and  manganese 
dioxide  are  used?     (c)  What  determines  which  is  the  cheaper  process? 

3.  The  labels  on  cans  of  bleaching  powder  usually  contain  a  statement 
as  to  the  available  chlorine  present  in  the  powder.     Available  chlorine  is 
that  which  is  set  free  when  bleaching  powder  is  treated  with  an  acid.     If  the 
amount  of  this  chlorine  is  stated  as  35  per  cent,  what  is  the  percentage  purity 
of  the  powder? 

4.  How  could  you  distinguish  by  chemical  means  the  folio  wing:  (a)  CaOCl2 
and  CaCl2,   (b)  NaCl  and  Na2SO4,    (c)  MnO2  and  C? 

5.  Hydrogen  is  a  very  light  gas  and  chlorine  is  very  heavy.     If  a  cylinder 
containing  the  former  is  inverted  and  placed  above  one  containing  chlorine, 
the  two  gases  will  slowly  mix  and,  after  a  time,  they  will  be  uniformly  dis- 
tributed.    Can  you  suggest  any  hypothesis  to  explain  this  fact? 

6.  If   100  liters    of  liquid    chlorine  were  allowed  to  change  to   a  gas 
what  volume  would  it  occupy  at  0°  and  760  mm.?     (b)  If  the  gas  were  mixed 
with  air  so  that  it  was  present  in  the  proportion  of  1  part  to  1,000,000  of  air 
what  would  be  the  volume  of  the  mixture?     (c)  Assuming  that  the  mixture 
formed  a  layer  10  meters  high  what  area  would  it  cover? 

7.  (a)  Into  what  compound  did  the  large  quantities  of  chlorine  liberated 
in  the  recent  war  probably  change?     (6)  What  becomes  of  the  chlorine  that 
escapes  in  a  chemical  laboratory?     (c)   Why  is  chlorine-water  kept  in  bottles 
of  brown  glass? 

8.  Can  you  think  of  any  use  in  warfare  that  might  be  made  of  the  fact 
that  chlorine  reacts  with  certain  hydrocarbons  to  form  carbon  and  hydro- 
chloric acid?     The  reaction  in  the  case  of  acetylene  is  practically  instantaneous. 

9.  Chlorine  is  both  a  disinfectant  and  a  deodorizer.     Why? 

10.  How  could  you  find  out  if  a  sample  of  hydrogen  chloride  contained  a 
small  amount  of  chlorine,  if  there  were  not  enough  of  the  latter  present  to 
be  distinguished  by  its  color? 

11.  Which  of  the  substances  indicated  by  the  following  formulas  would 


122  INORGANIC  CHEMISTRY  FOR  COLLEGES 

you  expect  to  be  inflammable?  (a)  CO2,  (6)  CC14,  (c)  CHC13,  (d)  CH4, 
(e)  PH3,  (/)  P205,  (<7)  C6H6,  (h)  C2H6O,  (i)  NO2. 

12.  Convert  the  following  into  balanced  equations:   (a)  C^He  +  O«  — *  CO2 
+  H20;      (6)  SbCl3  +  H2O  ->  Sb2O3  +  HC1;      (c)  KC1  +  H2SO4  -»  HC1  + 
+  K2SO4;  (d)  CH4  +  Br2  ->  CHBr3  +  HBr;  (c)  C2H6O  +  O2  ->  CO2  +  H2O. 

13.  The  element  bromine  forms  a  compound  of  the  formula  HBr:    (a) 

What  is  the  valence  of  the  element?    (6)  Write  the  formulas  of  the  com- 

i      ii     in   in  n 

pounds  of  bromine  with  Na,  Ca,  Al,  Sb,  and  Zn.  (c)  What  is  the  valence 
of  sulphur  in  the  compound  ZnS?  (d)  Write  the  formulas  of  the  compounds 
of  sulphur  and  Na,  Ca,  and  Sb.  (e)  What  is  the  valence  of  the  SO3  radical 

in  H2SO3?     (/)  Write  the  formulas  of  compounds  containing  this  radical 

i  n 

and  K  and  Ca.     (g)  The  formula  of  nitric  acid  is  HNO3,  write  the  formulas 

i     n  in 

of  the  compounds  containing  the  NO3  radical  and  K,  Cu,  and  Al. 

14.  At  0°  100  volumes  of  water  dissolve  461  volumes  of  chlorine.     Cal- 
culate the  percentage  of  chlorine  by  weight  in  the  mixture. 

15.  Calculate  the  weight  of  1  liter  of  chlorine,  C12,  at  0°  and  760  mm.  from 
its  atomic  weight. 

16.  If  2  liters  of  chlorine  are  dissolved  in  1  liter  of  water  and  the  mixture 
exposed  to  the  sunlight  so  that  the  oxygen  formed  is  collected,  what  volume 
of  the  latter  would  be  obtained? 


CHAPTER  XI 
HYDROCHLORIC  ACID.     DOUBLE  DECOMPOSITION 

136.  The  study  of  hydrochloric  acid  will  furnish  an  opportunity 
to  learn  a  great  deal  about  a  very  important  class  of  substances, 
called  acids,  which  possess  many  properties  in  common.     We  come 
in  contact  often  in  daily  life  with  compounds  of  this  class.     Acids 
are  formed  when  many  organic  substances  undergo  decay  through 
the  influence  of  the  bacteria  present  in  the  air;  the  juices  of  fruits 
and  vegetables  ferment,  and  gases  and  acids  are  formed;  and  milk 
turns  sour. 

Hydrochloric  acid  is  of  particular  interest  because  it  plays  a 
very  important  part  in  the  processes  that  take  place  in  the  animal 
body.  The  digestion  of  certain  foods  in  the  stomach  is  brought 
about  through  the  action  of  hydrochloric  acid  and  a  substance 
called  pepsin,  both  of  which  are  present  in  gastric  juice.  Pepsin, 
which  can  be  obtained  for  experimental  purposes  from  the  lining 
of  a  pig's  stomach,  does  not,  alone,  bring  about  the  digestion 
of  foods,  but  in  the  presence  of  hydrochloric  acid  becomes  very 
active;  the  latter  acts  probably  as  a  catalytic  agent.  As  the 
hydrochloric  acid  enters  into  combination  with  other  substances 
and  is  excreted  from  the  body,  some  compound  which  will  serve 
as  a  source  of  this  acid  must  be  taken  with  the  food.  Chlorides 
are  present  in  many  foods;  as  these,  however,  do  not  generally 
furnish  an  adequate  supply  of  the  acid,  common  salt,  sodium 
chloride,  is  an  important  constituent  of  the  diet.  Animals  have 
been  known  to  risk  death  in  a  desire  to  get  salt.  Hydrochloric 
acid  is  frequently  prescribed  by  physicians  in  those  cases  of  dys- 
pepsia which  are  supposed  to  be  due  to  the  lack  of  the  acid  in  the 
gastric  juice. 

137.  Historical. — Hydrochloric  acid  was  known  to  the  early 
alchemists,  who  probably  prepared  it  by  heating  a  mixture  of  salt 
and  green  vitriol   (FeSO4,7H2O).     Glauber,   one    of    the    most 

123 


124  INORGANIC  CHEMISTRY  FOR  COLLEGES 

famous  of  the  alchemists,  described  in  1648  the  preparation  of  the 
acid  from  salt  and  sulphuric  acid.  The  other  product  formed  in 
the  reaction,  sodium  sulphate,  Na2S04,10H2O,  was  for  a  long 
time  known  as  Glauber's  salt.  Priestley  was  the  first  to  obtain 
the  acid  free  from  water,  and  collected  the  gas  over  mercury. 

138.  Preparation   of   Hydrochloric   Acid. — The   acid    can   be 
made  by  burning  hydrogen  in  chlorine — a  reaction  which  has  been 
discussed  in  some  detail  (120).     The  equation  for  the  reaction  is 
repeated  here: 

H2  +  C12  =  2HC1 

This  is  not  ordinarily  a  practical  method  because  it  involves  the 
previous  preparation  of  both  hydrogen  and  chlorine.  It  is  used, 
however,  to  manufacture  hydrochloric  acid  when  a  very  pure  acid 
is  required,  and  when  hydrogen  and  chlorine  are  available.  This 
is  the  case  when  sodium  hydroxide  is  made  by  the  electrolysis  of  a 
solution  of  sodium  chloride. 

The  method  used  in  the  laboratory  and  on  the  large  scale  is 
based  on  the  reaction  between  salt  and  concentrated  sulphuric 
acid.  When  the  two  substances  are  heated  together  at  a  high 
temperature  a  reaction  indicated  by  the  following  equation  takes 
place : 

2NaCl  +  H2SO4  =  Na2SO4  +  2HC1 

At  lower  temperatures — those  used  in  the  laboratory — the  sub- 
stances interact  in  different  proportions: 

NaCl  +  H2SO4  =  NaHS04  +  HC1 

The  products  of  the  reaction  in  the  first  case  are  sodium  sulphate 
and  hydrochloric  acid;  in  the  second  case  sodium  hydrogen 
sulphate  is  formed.  The  acid  escapes  as  a  gas  and  can  be  col- 
lected by  the  upward  displacement  of  air,  since  it  is  heavier  than 
air,  or  by  conducting  it  into  water  in  which  it  dissolves. 

139.  Physical   Properties    of   Hydrochloric   Acid. — The   acid 
is  a  colorless  gas,  which  is  heavier  than  air;    1  liter  at  0°  and 
760  mm.  pressure  weighs  1.641  grams.     One  liter  of  water  at  0° 
dissolves  525  liters  of  the  gas. 

In  order  to  distinguish  between  the  pure  compound  and  its 
solution  in  water,  the  former  is  sometimes  called  hydrogen  chloride. 


HYDROCHLORIC    ACID.     DOUBLE  DECOMPOSITION       125 


The  marked  tendency  of  hydrogen  chloride  to  dissolve  in  water 
can  be  shown  by  a  striking  experiment.  A  large  glass  globe  is 
filled  with  the  dry  gas  and  then  closed  by  a  stopper  carrying  a 
long  glass  tube  and  a  medicine  dropper  filled  with  water  and  so 
placed  that  by  means  of  it  a  small  amount  of  water  can  be  forced 
into  the  globe  (see  Fig.  15).  The  apparatus  is  then  inverted 
and  supported  over  a  reservoir  con- 
taining water  colored  by  blue  litmus. 
When  water  is  forced  into  the  globe  by 
pressing  the  rubber  end  of  the  medicine 
dropper,  the  gas  dissolves  in  it;  this 
leaves  a  partial  vacuum,  and  the  water 
in  the  reservoir  rushes  up  to  take  the 
place  of  the  dissolved  gas;  as  it  enters, 
the  blue  litmus  is  changed  to  red  by  the 
acid. 

140.  Chemical  Conduct  of  Hydro- 
chloric Acid. — When  the  gas  is  dry  it 
is  very  inactive,  but  a  small  amount  of 
water  serves  as  a  catalytic  agent  in 
helping  to  bring  about  reaction  be- 
tween it  and  other  substances.  Dry 
hydrochloric  acid  does  not  react  with 
zinc,  even  when  the  gas  is  dissolved 
in  a  liquid  such  as  benzene.  The 

addition  of  a  small  amount  of  water,  however,  causes  a  reaction  to 
begin,  and  hydrogen  is  evolved.  We  have  learned  that  acids 
change  the  color  of  litmus  paper  from  blue  to  red.  If  a  dry  piece 
of  blue  litmus  paper  is  put  into  the  dry  gas,  it  is  scarcely  affected; 
if  the  paper  is  first  moistened  it  turns  red  immediately. 

Hydrochloric  acid  as  ordinarily  prepared  contains  enough  water 
to  catalyze  the  reactions  between  it  and  other  substances.  Under 
these  circumstances  it  reacts  with  ammonia.  This  can  be  shown 
by  a  striking  experiment.  A  large  cylinder,  the  open  end  of  which 
is  carefully  ground  flat  and  covered  with  vaseline,  is  filled  with  the 
gas  by  upward  displacement  of  air.  It  is  then  closed  by  placing 
over  it  a  ground- glass  plate,  the  vaseline  serving  to  make  the 
covering  air-tight.  A  similar  cylinder  of  ammonia,  which  is  also  a 
colorless  gas  having  the  formula  NHs,  is  prepared  and  covered. 


FIG.  15. 


126  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  cylinders  are  now  placed  mouth  to  mouth,  the  one  containing 
ammonia  on  top.  The  two  glass  plates  are  next  quickly  with- 
drawn and  the  open  ends  of  the  cylinders  brought  into  close  con- 
tact. A  dense  white  cloud  appears  at  once,  formed  as  the  result 
of  the  combination  of  the  two  invisible  gases;  the  substance  is 
ammonium  chloride,  which  is  often  called  sal  ammoniac,  and  is 
used  in  certain  kinds  of  electric  batteries.  The  reaction  is  ex- 
pressed by  the  following  equation: 

HC1  +  NH3  =  NH4C1 

If  an  attempt  is  made  to  pull  apart  the  cylinders,  it  will  be  found 
that  considerable  force  must  be  exerted.  This  is  because  the 
gases  have  disappeared,  and  in  separating  the  cylinders  we  must 
overcome  the  pressure  of  the  atmosphere  on  them,  which  is  not 
appreciated  as  long  as  they  were  filled  with  gas.  The  cylinders 
can  be  separated  easily  by  sliding  one  over  the  other,  since  in 
doing  this  we  are  not  working  against  the  air-pressure.  When  a 
small  opening  is  made  in  this  way  the  air  rushes  in  to  fill  the 
vacuum. 

141.  When  a  jar  of  hydrochloric  acid  is  opened  to  the  air  a 
white  cloud  is  observed  at  the  mouth  of  the  jar.     The  cloud  is 
denser  if  the  breath  is  blown  through  the  gas.     Water  can  exist 
in  the  air  in  the  form  of  an  invisible  gas,  called  water-vapor. 
When  certain  substances  which  are  very  soluble  in  water  are 
brought  in  contact  with  moist  air,  they  absorb  water  from  the 
latter,  and,  finally,  a  solution  of  the  compound  is  formed.     The 
cloud  produced  from  hydrochloric  acid  consists  of  minute  drops 
of  a  solution  of  the  gas  in  water.     When  a  gas  behaves  in  this 
way,  it  is  said  to  fume  in  the  air.     Solids  which  take  up  moisture 
from  air,  become  damp,  and  finally  pass  into  solution  in  the 
water  absorbed,  are  said  to  be  deliquescent.     Calcium  chloride, 
Cadi,  is  an  example  of  such  a  salt. 

142.  Properties  of  Aqueous  Solutions  of  Hydrochloric  Acid.— 
When  hydrogen  chloride  is  dissolved  in  water,  the  solution  exhibits 
many  important  properties  which  are  not  shown  by  the  anhy- 
drous compound.     The  word  anhydrous  means  free  from  water; 
it  is  a  very  convenient  adjective  and  is  often  used  in  chemistry. 

The  solution  of  hydrochloric  acid  in  water,  like  the  solutions 
of  other  acids,  has  a  sour  taste.  It  turns  blue  litmus  paper  red, 
and  effects  a  change  in  color  in  other  substances.  The  dye  called 


HYDROCHLORIC  ACID.     DOUBLE  DECOMPOSITION       127 

methyl  orange  is  yellow ;  it  is  changed  to  a  red  compound  when 
treated  with  an  acid.  The  salts  of  phenolphthalein  are  red,  but 
are  converted  into  a  colorless  compound  by  acids.  Substances 
which  behave  in  this  way  are  called  indicators;  they  are  much 
used  in  chemistry  to  detect  the  presence  of  acids. 

Hydrochloric  acid  reacts  with  the  more  active  metals  and,  as 
a  result,  hydrogen  and  a  chloride  are  formed.  A  number  of 
such  reactions  have  already  been  discussed  (46).  Two  typical 
cases  are  represented  by  the  following  equations: 

Zn  +  2HC1  =  ZnCl2  +  H2 
2A1  +  6HC1  =  2A1C13  +  3H2 
Hydrochloric  acid  reacts  with  the  oxides  of  metals: 
Na20  +  2HC1  =  2NaCl  +  H20 
FeO  +  2HC1  =  FeCl2  +  H2O 
A1203  +  6HC1  =  2A1C13  -h  3H20 
The  acid  reacts  also  with  hydroxides : 

NaOH  +  HC1  =  NaCl  +  H20 
Zn(OH)2  +  2HC1  =  ZnCl2  +  2H2O 
A1(OH)3  +  3HC1  =  A1C13  +  3H20 

All  acids  enter  into  reactions  similar  to  those  illustrated  by  means 
of  hydrochloric  acid  in  the  three  series  of  equations  just  given. 

143.  Commercial  hydrochloric  acid,  which  is  often  called  muri- 
atic acid,  possesses  a  yellow  color  which  is  caused  by  the  presence 
of  impurities  in  it.  It  usually  contains  small  quantities  of  sulphuric 
acid,  chlorine,  iron  chloride,  and  arsenic.  It  contains  about  40  per 
cent  by  weight  of  hydrogen  chloride,  the  rest  being  chiefly  water. 
It  is  sold  in  large  glass  vessels  called  carboys;  each  vessel  is  enclosed 
in  a  wooden  box  and  is  packed  in  straw.  Carboys  for  shipment 
of  acid  generally  hold  about  12  gallons.  A  purer  quality  of  the 
acid  can  be  purchased:  this  is  sold  under  the  designation  C.P., 
the  letters  being  an  abbreviation  of  the  words  chemically  pure. 
So-called  C.P.  chemicals  are  supposed  to  contain  not  more  than 
traces  of  impurities,  the  amounts  of  these  present  not  being  enough 


128  INORGANIC  CHEMISTRY  FOR  COLLEGES 

to  interfere  with  the  use  of  the  chemical  for  all  ordinary  purposes. 
It  takes  a  very  long  time  to  obtain  any  substance  in  such  a  pure 
condition  that  the  presence  of  other  substances  in  minute  amounts 
cannot  be  detected  in  it  by  a  skilled  chemist.  Hydrochloric  acid, 
C.P.,  contains  40  per  cent  by  weight  of  HC1  and  has  the  specific 
gravity  1.2,  that  is,  it  is  1.2  times  as  heavy  as  an  equal  volume  of 
water.  Acid  of  this  strength  is  called  concentrated  hydrochloric 
acid;  it  is  made  by  passing  the  gas  into  water  as  long  as  it  is 
absorbed  at  ordinary  temperatures;  such  a  solution  is  said  to  be 
saturated. 

Dilute  hydrochloric  acid  can  be  made  of  any  strength  less 
than  this;  that  commonly  supplied  in  the  laboratory  contains 
about  20  per  cent  of  HC1.  What  happens  when  hydrochloric  acid 
is  heated  is  determined  by  the  proportion  of  hydrogen  chloride  it 
contains.  If  the  solution  is  concentrated,  the  gas  is  driven  off 
and  if  dilute,  water-vapor  first  escapes.  In  either  case  as  the  acid 
is  heated  a  point  is  finally  reached  when  the  product  that  distills 
over  boils  at  110°  at  760  mm.,  and  contains  20.24  per  cent  HC1, 
and  79.76  per  cent  EkO.  This  is  called  the  constant-boiling  mix- 
ture of  water  and  hydrochloric  acid. 

Hydrochloric  acid  is  used  in  the  preparation  of  chlorine  and  of 
chlorides,  in  cleaning  metals,  and  for  various  purposes  in  the 
chemical  laboratory. 

144.  Chlorides. — Chlorides  can  be  formed  by  the  action  of 
chlorine  on  the  various  elements;    a  number  of  examples  have 
already  been  noted  (121).     In  the  case  of  the  metals,  chlorides 
can  be  prepared  by  means  of  the  reactions  which  were  illustrated 
by  the  equations  given  in  the  last  section.     Many  metals  which 
do  not  react  with  hydrochloric  acid  form  oxides  which  are  converted 
into  chlorides  by  the  acid.     For  example,  copper  and  hydrochloric 
acid  do  not  react,  but  copper  oxide  dissolves  readily  in  the  acid: 

CuO  +  2HC1  =  CuCl2  +  H2O 

The  uses  to  which  chlorides  are  put  will  be  described  in  connection 
with  the  consideration  of  the  metals  which  they  contain. 

145.  Test  for  Chlorides. — It  is  often  necessary  to  determine 
whether  any  particular  substance  is  a  chloride,  or  whether  a  mix- 
ture contains  a  chloride.     Such  problems  arise  in  what  is  known  as 
qualitative  analysis — a  branch  of  chemistry  which  has  to  do  with 


HYDROCHLORIC  ACID.    DOUBLE  DECOMPOSITION       129 

the  determination  of  the  presence  of  the  various  substances  in  com- 
mercial and  other  products.  We  turn  to  the  chemist  to  analyze  a 
supposed  poison,  a  sample  of  steel  or  brass,  or  a  patent  medicine; 
to  test  milk  and  drinking  water  for  impurities,  or  a  rock  for  gold. 
All  such  problems  are  solved  by  qualitative  analysis.  The  pres- 
ence of  any  substance  in  a  mixture  is  established  by  separating 
and  identifying  it,  or  by  converting  it  into  another  substance  the 
properties  of  which  are  known.  The  method  used  can  be  illus- 
trated clearly  by  a  consideration  of  the  way  in  which  the  presence 
of  a  chloride  is  determined.  In  this  case  it  is  not  necessary  to 
separate  the  chloride  itself,  provided  all  we  want  to  know  is  whether 
a  chloride  is  present.  The  substance  to  be  tested  is  dissolved  in 
water;  if  it  is  insoluble  in  water  another  liquid  is  used.  A  solu- 
tion of  silver  nitrate  is  next  added.  If  a  chloride  is  present  a  white 
cloud  appears  or  a  white  solid  separates  and  settles  to  the  bottom 
of  the  tube.  When  an  insoluble  substance  separates  in  this  way 
on  mixing  two  solutions,  it  is  called  a  precipitate.  The  substance 
in  this  case  is  silver  chloride;  the  reaction  is  represented  by  the 
following  equation  when  the  chloride  present  is  sodium  chloride: 

NaCl  +  AgNO3  =  NaNO3  +  AgCl 

It  is  seen  that  the  metals  change  places,  and  we  get  as  the  result  of 
the  reaction  the  chloride  of  silver  and  the  nitrate  of  the  metal 
which  was  originally  in  combination  with  chlorine.  The  chlorides 
of  the  metals  in  general  act  in  this  way,  for  example: 

ZnCl2  +  2AgNO3  =  Zn(NO3)2  +  2AgCl 
A1C13  +  3AgN03  =  A1(NO3)3  +  3AgCl 

We  could  use  the  formation  of  the  white  precipitate  when  silver 
nitrate  is  added  to  the  solution,  as  a  test  for  chloride,  provided  no 
other  substances  produced  a  similar  result.  Other  substances, 
however,  do  give  a  white  precipitate  under  these  conditions,  and 
we  must  find  some  property  of  silver  chloride  other  than  its  color 
and  insolubility  in  water,  if  we  are  to  use  its  formation  to  show  the 
presence  of  chlorides.  Silver  chloride  is  insoluble  in  dilute  nitric 
acid,  whereas  the  white  precipitates  formed  when  substances  other 
than  chlorides  are  treated  with  silver  nitrate  are  soluble  in  this 
acid.  This  important  fact  makes  it  possible  to  use  the  reaction 


130  INORGANIC  CHEMISTRY  FOR  COLLEGES 

with  silver  nitrate  as  a  test  for  a  chloride;  we  add  dilute  nitric 
acid  to  the  solution  before  adding  silver  nitrate,  and  if  a  white 
precipitate  is  formed  we  conclude  that  a  chloride  is  present. 
This  method  of  detecting  a  chloride  in  the  presence  of  other  sub- 
stances is  applicable  in  the  case  of  practically  all  commonly  occur- 
ring compounds ;  it  is  used  as  a  test  for  chlorides. 

146.  When  we  wish  to  test  for  a  substance  or  one  of  its  con- 
stituents, we  convert  it  into  a  new  substance  possessing  charac- 
teristic properties  which  can  be  readily  recognized,  and  the  sum 
of  which  are  not  possessed  by  any  other  substance  that  could  be 
formed  under  the  conditions  used  in  the  test.     We  use  at  times  the 
production  of  an  odor,  a  color,  or  an  insoluble  substance,  the  forma- 
tion of  the  latter  being  accomplished  if  possible.     In  learning  a 
test  the  student  should  take  care  to  understand  what  are  the 
characteristic  properties  of  the  substances  upon  the  formation  of 
which  the  test  is  based.     The  test  for  a  chloride  is  based  on  the 
fact  that  silver  chloride  is  a  white  solid  which  is  insoluble  in  nitric 
acid.     Since  silver  bromide  is  insoluble  in  nitric  acid  and  possesses 
a  faint  yellow  color,  it  might  be  mistaken  by  the  beginner  for  silver 
chloride;  to  avoid  this,  a  solution  of  ammonia  is  added  to  the  pre- 
cipitate;  silver  chloride  dissolves  readily,  whereas  silver  bromide 
requires  a  large  amount  of  the  solution  to  dissolve  it.     It  should 
be  noted  here,  also,  that  silver  chloride  when  exposed  to  sunlight 
soon  turns  in  color  to  a  light  violet  shade;  silver  bromide  does  not 
do  this.     In  order  to  be  sure  that  the  conclusion  drawn  from  a  test 
is  correct,  it  is  advisable  for  the  beginner  to  apply  independent 
tests,  if  possible,  to  confirm  his  conclusion.     A  test  made  for  this 
purpose  is  called  a  confirmatory  test.     Accumulation  of  evidence 
in  support  of  a  statement  makes  it  more  worthy  of  belief;    we 
cannot  be  too  careful  in  drawing  our  conclusions. 

147.  Types  of  Chemical  Reactions. — A  large  number  of  chem- 
ical reactions  have  been  discussed  up  to  this  point.     It  is  advisable 
to  reexamine  some  of  them  and  see  how  they  can  be  put  into  a  few 
simple  classes.     One  of  the  chief  aims  of  science  is  to  systematize 
knowledge,  and  as  has  been  pointed  out,  the  student  should  con- 
stantly attempt  to  correlate  the  new  facts  as  they  are  presented, 
that  is,  to  see  their  relation  to  facts  already  mastered. 

The  simplest  type  of  reaction  which  has  been  studied  is  that 
which  occurs  when  two  substances  unite  to  form  a  third  substance; 


HYDROCHLORIC  ACID.    DOUBLE  DECOMPOSITION       131 

this  is  called  combination  and  is  illustrated  by  many  equations 
already  given,  for  example: 

2Cu  +  O2  =  2CuO 

The  reverse  of  combination  is  decomposition: 
2HgO  =  2Hg  +  02 
Displacement  is  a  third  type  of  reaction: 

Zn  +  H2SO4  =  ZnS04  +  H2 

In  the  above  case  hydrogen  has  been  substituted  for  zinc  in  H2S04. 
148.  A  fourth  type  is  illustrated  by  such  a  reaction  as  the 
following: 

NaCl  +  AgN03  =  AgCl  +  NaN03 

In  this  case  since  two  substances  interact  and  form  two  new  sub- 
stances, the  reaction  is  called  one  of  double  decomposition  or  meta- 
thesis. A  reaction  is  put  into  this  class  only  when  the  two  sub- 
stances interact  in  a  particular  way,  which  will  be  explained  in 
detail,  since  reactions  of  double  decomposition  are  examples,  per- 
haps, of  the  most  important  type  that  the  student  will  meet. 

The  significance  of  the  expression  acid  radical  has  been 
explained;  it  is  the  group  of  atoms  other  than  hydrogen  present 
in  the  acids.  For  example,  the  radicals  present  in  the  acids 
having  the  formulas  HC1,  HNO3,  H2SO4,  H2SO3,  H3PO4,  are  Cl, 
N03,  SO4,  SO3,  and  PO4,  respectively.  Compounds  made  up  of 
metallic  atoms  and  acid  radicals  are  called  salts,  for  example, 
AgNO3,  Na2SO4,  etc.  It  will  also  be  recalled  that  bases  are  com- 
pounds which  contain  metallic  atoms  in  combination  with  one  or 
more  hydroxyl  groups,  such  as  NaOH,  Zn(OH)2,  etc.;  the  hydroxyl 
group  has  the  valence  1.  Reactions  of  double  decomposition 
take  place  between  pairs  of  compounds  belonging  to  these  classes. 
With  these  facts  before  us  we  can  examine  fully  the  reactions  of 
this  type  which  have  been  studied,  and  become  prepared  to  under- 
stand those  that  are  to  follow.  An  examination  of  the  equation 
last  given, 

NaCl  +  AgN03  =  AgCl  +  NaN03, 

brings  out  the  fact  that  the  reaction  is  between  two  salts,  and  that 
the  two  products  of  the  reaction  are  formed  as  the  result  of  an 


132  INORGANIC  CHEMISTRY  FOR  COLLEGES 

exchange  of  the  metallic  atoms;  sodium  takes  the  place  of  silver, 
and  silver  that  of  sodium.  A  much  better  name  than  double 
decomposition  for  this  class  of  reaction  would  be  double  sub- 
stitution. 

We  have  seen  that  metallic  atoms  can  take  the  place  of 
hydrogen  in  acids;  an  example  of  this  kind  of  reaction  is  given 
above  as  an  example  of  substitution.  Such  an  exchange  is  possible 
in  double  decomposition  reactions;  the  formation  of  hydrogen 
chloride  from  salt  and  sulphuric  acid  is  an  example: 

2NaCl  +  H2SO4  =  Na2S04  +  2HC1 

In  this  case  a  salt  and  an  acid  enter  into  double  decomposition.  It 
is  seen  from  the  two  examples  given  that  a  double  decomposition 
takes  place  between  two  substances  when  their  hydrogen  or  metallic 
atoms  mutually  replace  each  other.  A  double  decomposition 
occurs  also  when  only  single  atoms  exchange  places,  for  example, 

NaCl  +  H2S04  =  NaHS04  +  HC1 

In  this  case  but  1  hydrogen  of  the  sulphuric  acid  changed  place 
with  1  sodium  atom.  Cases  of  this  kind  are  not  so  common  as 
those  in  which  all  the  atoms  united  with  the  acid  radical  enter  into 
the  exchange. 

A  few  double  decompositions  will  be  considered  in  detail  to 
illustrate  how  reactions  of  this  type  can  be  written  if  we  apply 
the  statement  just  given.  Suppose  it  is  desired  to  represent  a 
double  decomposition  between  the  compounds  sodium  chloride 
and  silver  sulphate.  We  first  set  down  the  formulas  of  these 
compounds  and  in  order  to  indicate  how  the  molecules  break  apart 
draw  dotted  lines  between  the  metallic  or  hydrogen  atoms  and  the 
acid  radicals;  we  also  indicate  their  valencies  thus: 

ill  i      i     II 

Na  |  Cl  +  Ag2  |  S04 

We  then  change  the  positions  of  the  metallic  atoms,  neglecting  the 
subscript  numbers  representing  the  number  of  atoms  involved  in 
everything  except  the  acid  radicals,  thus : 

I     i     I  I      !      II  ill  I     !     II 

Na  !  Cl  +  Ag2  I  S04  -»  Ag  |  C1  +  Na  !  S04 

i  I  i  i 


HYDROCHLORIC  ACID.    DOUBLE  DECOMPOSITION       133 

We  next  add  subscript  numbers  which  are  necessary  to  satisfy  the 
valence  requirement.  In  the  above  case  a  2  must  be  placed  to  the 
right  and  below  the  symbol  for  sodium  in  sodium  sulphate  in  order 
that  the  salt  may  have  the  correct  formula.  Finally  the  equation 
is  balanced ;  this  is  done  by  taking  two  molecules  of  NaCl  and  two 
of  AgCl.  The  equation  becomes  as  a  result: 

2NaCl  +  Ag2SO4  =  2AgCl  +  Na2SO4 

Another  double  decomposition  will  be  written  in  this  way  between 
aluminium  chloride  and  sulphuric  acid.  First  represent  the  parts 
which  will  interchange  and  the  result  of  the  interchange,  affixing 
the  numbers  representing  the  valence  of  each, 

in  i    i          i    j    ii          i   i    i        in  ;    ii 
Al  !  C13  +  H2  !  SO4  -*  H  !  Cl  +  Al  I  S04 
i  i  !  ! 

Next  add  the  subscript  numbers  to  satisfy  the  valence  require- 
ments : 

in  i    i          i     i    ii          iii         in  !     ii 

Al  |  C13  +  H2  I  SO4  ->  H  !  Cl  +  A12  !  (SO4)3 

|  j  I  I 

Finally  balance  the  equation: 

2A1C13  +  3H2S04  =  6HC1  +  A12(S04)3 

All  these  operations  can  be  done  in  successive  steps  without  writing 
the  equation  three  times  as  was  done  here  to  make  each  step  clear. 
Double  decompositions  which  involve  bases  are  written  in  the 
same  way: 

I     i        I  III  I      I  III    !       I  I       i     I 

Na  |  OH  +  Al  j  Cla  -»  Al  |  OH  +  Na  I  Cl 

ill  i 

3NaOH  +  A1C13  =  A1(OH)3  +  3NaCl 
In  double  decompositions  between  acids  and  bases  water  is  formed : 

Na  j  OH  +  H2  |  S04  =  Na  |  SO4  +  H  |  OH 

2NaOH  +  H2SO4  =  Na2S04+  2H2O 

From  the  above  considerations  we  see  that  double  decompositions 
can  be  written  between  the  following  pairs  of  substances:  Acids 


134  INORGANIC  CHEMISTRY  FOR  COLLEGES 

and  salts,  acids  and  bases,  bases  and  salts,  and  salts  and  salts.  It 
is  right  to  ask  at  this  point,  do  all  the  reactions  written  in  this  more 
or  less  mechanical  way  represent  chemical  changes  that  actually 
take  place?  It  has  been  already  emphasized  that  it  is  impossible 
to  arrive  at  a  chemical  formula  or  reaction  by  a  simple  mathe- 
matical juggling  of  symbols.  We  shall  see  later  that  all  the  chem- 
ical changes  indicated  by  double  decomposition  reactions  between 
the  pairs  of  substances  listed,  do  take  place  to  a  greater  or  less 
extent.  Only  when  the  reaction  is  practically  complete  in  the 
manner  indicated  by  the  equations  do  we  call  them  reactions  of 
double  decomposition. 

149.  A  study  of  many  reactions  of  this  type  has  brought  out 
the  important  generalization  that  if  one  of  the  products  formed  can 
escape  in  any  way,  the  double  decomposition  takes  place.  When, 
for  example,  a  gas  is  formed,  it  leaves  the  vessel  in  which  the  reac- 
tion is  taking  place.  The  reaction  also  takes  place  if  one  of  the 
products  is  insoluble;  it  precipitates  and  is  no  longer  in  solution 
where  it  can  react;  it  leaves  the  field  of  action.  We  can  readily 
see  why  this  removal  of  one  of  the  products  affects  the  course  of  a 
reaction.  Double  decompositions  belong  to  an  important  class 
of  reactions  known  as  reversible  reactions.  In  reactions  of  this 
kind  the  substances  indicated  on  the  left-hand  side  of  the  equation 
can  interact  to  form  those  indicated  on  the  right-hand  side,  and 
those  on  the  right  can  form  those  on  the  left.  We  express  this  in 
the  equation  for  the  reaction  by  replacing  the  sign  of  equality  by 
another  symbol,  thus : 

NaCl  +  H2SO4  <=±  NaHSO4  +  HC1 

Written  in  this  way,  the  equation  indicates  that  sodium  chloride 
and  sulphuric  acid  can  react  to  form  sodium  hydrogen  sulphate 
and  hydrochloric  acid,  and,  also,  that  sodium  hydrogen  sulphate 
and  hydrochloric  acid  can  react  to  form  sodium  chloride  and  sul- 
phuric acid.  It  might  well  be  asked,  under  what  circumstances 
does  the  reaction  take  place  in  one  direction  only.  It  is  evident 
that  if  the  hydrochloric  acid  is  taken  away  as  soon  as  it  is  formed, 
it  cannot  react  with  the  sodium  hydrogen  sulphate  produced  with 
it  to  form  sodium  chloride  again.  We  are  now  in  a  position  to 
understand  under  what  circumstances  reversible  reactions  of  this 
kind  proceed  to  completion  in  one  direction,  and  become,  thus, 


HYDROCHLORIC  ACID.     DOUBLE  DECOMPOSITION       135 

reactions  of  double  decomposition.  This  occurs  when  one  of  the 
products  is  removed,  the  removal  taking  place  of  itself  if  a  gas 
which  escapes  is  formed,  or  if  an  insoluble  substance  is  produced 
which  separates  as  a  precipitate.  With  this  knowledge  we  can 
examine  any  reaction  written  between  the  pairs  of  substances  listed 
above,  and  state  whether  they  represent  double  decomposition 
reactions.  In  order  to  do  this  it  is  necessary  to  know  whether 
either  of  the  substances  produced  is  a  gas  which  escapes  or  an 
insoluble  compound. 

Double  decomposition  also  takes  place  when  water  is  one  of  the 
products  of  the  reaction;  the  reason  for  this  will  be  explained  fully 
later.  In  brief,  then,  reactions  of  double  decomposition  take  place 
between  acids,  bases,  and  salts  when  one  of  the  products  is  a  gas, 
an  insoluble  substance,  or  water.  This  is  a  generalization  of  the 
greatest  importance,  because  a  large  proportion  of  the  reactions 
to  be  met  with  belong  to  this  type, 

EXERCISES 

1.  Calculate  from  the  formula  of  hydrochloric  acid  the  weight  of  1  liter 
of  the  gas  at  0°  and  760  mm. 

2.  (a)  What  is  meant  by  the  expression  "water  seeks  its  own  level"? 
(6)  What  causes  the  water  to  rise  into  the  globe  in  the  experiment  described 
in  section  139?     (c)  How  would  you  tell  after  the  experiment  how  much  air 
was  mixed  with  the  hydrogen  chloride  in  the  globe? 

3.  (a)  Why  does  concentrated  hydrochloric  acid  fume  in  the  air?     (6) 
Why  does  it  fume  more  strongly  if  the  breath  is  blown  across  the  open  mouth 
of  the  bottle? 

4.  Write  equations  for  the  reactions  which  take  place  between  hydro- 
chloric acid  and  the  following:     (a)  ZnO,     (6)  K2O,     (c)  Fe,     (d)  Fe2O3, 
(e)  Sn,   (/)  Mg,     (g)  Fe(OH)2,    (h)  A1(OH)3. 

5.  (a)  Write  equations  for  three  different  reactions  by  which  ZnClc  can 
be  made.     (6)  Can  CuCl2  be  made  by  similar  reactions? 

6.  Write  equations  for  reactions  involved  when  the  following  are  tested 
to  determine  whether  they  are  chlorides :  HC1,  CaCl2,  FeCl3. 

7.  (a)  Is  silver  nitrate  necessary  if  it  is  desired  to  test  for  a  chloride? 
(6)  If  not,  what  other  salt  could  be  used?     (c)  Write  equations  for  the  reac- 
tions which  would  take  place  if  this  salt  were  used  in  testing  solutions  of 
NaCl,  ZnCl2,  and  A1C13. 

8.  (a)  How  much  hydrochloric  acid,  HC1,  is  required  to  dissolve  10  grams 
of  zinc?     (b)  If  concentrated  hydrochloric  acid  which  contains  40  per  cent 
HC1  is  used,  how  much  of  it  is  necessary  to  dissolve  the  zinc?     (c)  Acid  of 
this  strength  has  the  specific  gravity  1.2.    What  volume  of  the  acid  is  required? 


136  INORGANIC  CHEMISTRY  FOR  COLLEGES 

9.  How  many  grams  of  (a)  concentrated  hydrochloric  acid  and  (b)  the 
constant-boiling  mixture  of  HC1  and  water  must  be  taken  to  obtain  1  gram- 
molecular-weight  of  HC1?  (c)  With  what  weight  of  sodium  hydroxide  will 
this  amount  of  acid  react?  (d)  If  this  weight  of  acid  is  mixed  with  enough 
water  to  make  the  volume  of  the  mixture  1  liter,  how  much  sodium  hydroxide 
is  present  in  a  solution  which  reacts  with  50  c.c.  of  the  solution  of  the  acid? 

10.  Using  the  method  described  in  the  text,  write  reactions  of  double  decom- 
position between  the  following  pairs  of  compounds:    (a)  NaOH  and  HC1, 
(6)  KOH  and  H2SO4,     (c)  NaOH  and  H3PO4,    (d)  CaCl2  and  Na2SO4,    (e) 
A1C13  and  KOH,  (/)  A1C13  and  H2SO4,   (g)  CaCl2  and  Na2PO4. 

11.  What  weight  of  silver  chloride,  AgCl,  is  formed  when  10  grams  of  silver 
nitrate,  AgNOs,  are  treated  with  hydrochloric  acid? 

12.  A  piece  of  a  silver  dime  weighing  0.200  gram  was  dissolved  in  nitric 
acid  and  the  solution  of  silver  nitrate,  AgNO>,  formed  was  treated  with  hydro- 
chloric acid.      The  silver  chloride  which  was  precipitated  was  dried  and 
found  to  weigh  0.239  gram,     (a)  How  much  silver  was  present  in  this  weight 
of  silver  chloride?     (b)  What  percentage  of  silver  did  the  dime  contain? 

13.  What  weight  of  sodium  chloride  is  required  to  furnish  enough  hydrogen 
chloride  to  fill  at  0°  and  760  mm.  a  flask  of  5  liters  capacity? 


CHAPTER  XII 
THE  ENERGY  FACTOR  IN  CHEMICAL  CHANGE 

150.  The  important  fact  has  already  been  emphasized  that  in 
all  chemical  changes  the  transformations  brought  about  in  mat- 
ter are  associated  with  changes  in  chemical  energy.      Energy 
manifests  itself  through  the  effect  it  produces  on  matter;    these 
effects  alone  appeal  to  our  senses,  but  we  must  study  their  cause 
if  we  are  to  get  at  a  fundamental  conception  of  matter  itself. 
Chemical  energy  cannot  be  directly  measured;   we  determine  the 
amount  of  this  form  of  energy  involved  in  a  change  by  transforming 
it  into  other  kinds  of  energy  which  can  be  measured.     It  has 
already  been  pointed  out  that  one  form  of  energy  can  be  changed 
into  another  form   (12).     When  this  fact  was   studied    quanti- 
tatively it  was  discovered  that  in  the  transformations  no  energy 
was  lost  or  gained.     A  number  of  investigators  announced  inde- 
pendently what  is  known  as  the  law  of  the  conservation  of  energy 
(J.  R.  Mayer  in  1842  and  Helmholtz  in  1847).     The  law  states 
that  within  a  system  undergoing  change  there  is  no  loss  or  gain 
in  energy.     In  certain  cases  one  kind  of  energy  can  be  com- 
pletely transformed  into  another  kind;  whereas  in  other  cases  the 
transformation   brings   about   the   appearance   of  two   kinds  of 
energy.     Electrical  energy  can  be  completely  converted  into  heat, 
but  when  an  attempt  is  made  to  change  it  into  light,  a  large  part 
of  the  energy  is  transformed  into  heat.     Changes  of  the  first  kind 
are  used  in  measuring  energy. 

151.  When  two  substances  enter  into  a  chemical  reaction, 
chemical  energy  is  changed  into  other  forms  of  energy.     It  is 
possible  to  carry  out  these  reactions  in  such  a  way  that  heat  is 
the  only  form  of  energy  produced.     Since  the  amount  of  heat 
can  be  measured,  the  total  change  in  energy  that  accompanies 
a  chemical  reaction  can  be  determined.     We  make  use  of  such 

137 


138  INORGANIC  CHEMISTRY  FOR  COLLEGES 

measurements,  for  example,  in  comparing  the  values  of  coal,  gas, 
alcohol,  etc.,  as  sources  of  heat  when  they  are  burned. 

In  comparing  the  chemical  behavior  of  two  substances,  for 
example,  the  behavior  of  iron  and  of  silver  toward  chlorine,  we  are 
concerned  with  the  relative  tendencies  of  the  two  metals  to  enter 
into  the  reaction  involved.  The  tendency  of  one  element  to 
react  with  another  to  form  a  compound  is  not  measured  by  the 
heat  produced  when  the  reaction  takes  place,  but  by  the  work 
necessary  to  reverse  the  chemical  change.  For  example,  in  the 
case  of  the  formation  of  silver  chloride  from  silver  and  chlorine 
the  tendency  for  the  compound  to  form  is  measured  by  the  work 
required  to  separate  one  gram-molecular-weight  of  the  chloride  into 
silver  and  chlorine  at  the  pressure  of  one  atmosphere.  It  is  evident 
that  the  greater  the  tendency  for  a  compound  to  form,  the  greater 
must  be  the  work  required  to  separate  it  into  its  constituents. 

The  work  involved  in  such  separations  can  be  measured  in 
several  ways.  One  of  these  involves  the  determination  of  the 
amount  of  electrical  energy  required  to  effect  the  decomposition 
of  the  compound  into  the  elements  from  which  it  was  formed. 
For  example,  we  can  measure  the  amount  of  electrical  energy 
required  to  separate  one  gram-molecular-weight  of  silver  chloride, 
into  silver  and  chlorine  at  the  pressure  of  one  atmosphere.  This 
work  is  taken  as  the  measure  of  the  tendency  of  silver  and 
chlorine  to  unite. 

152.  When  chemical  changes  take  place,  heat  is  produced  or 
disappears.     The  importance  of  this  fact  was  recognized  early, 
and  measurements  of  the  heat  change  in  many  chemical  reactions 
have  been  made.     Although  these  heat  values  are  not  a  measure 
of  the  tendencies  of  the  reactions  to  take  place,  they  are  helpful, 
however,   in  comparing  the  relative  tendencies  in  the  case  of 
similar  reactions.     It  has  been  found,   for  example,   that  the 
tendencies  of  the  several  metals  to  unit  with  chlorine  are  in  the 
same  order  as  the  values  of  the  heat  produced  when  the  chlorides 
of  these  metals  are  formed. 

153.  Measurement  of  Heat  Energy. — We  must  next  turn  our 
attention  to  the  way  in  which  heat  energy  is  measured ;  electrical 
energy  will  be  considered  later. 

We  are  familiar  with  the  thermometer  and  know  that  it  is  used 
to  measure  temperature;  it  tells  us  how  hot  a  thing  is  but  it  gives 


THE  ENERGY  FACTOR  IN  CHEMICAL  CHANGE 


139 


us  no  indication  of  how  much  heat  is  present.  If  we  were  to  test 
the  flame  of  a  burning  match  we  should  find  that  it  was  as  hot  as 
the  flame  from  a  log  of  wood ;  the  two  would  show  approximately 
the  same  temperature,  but  with  one  we  could  heat  a  room  and  with 
the  other  we  could  not ;  the  quantity  of  heat  is  different  in  the  two 
cases.  We  see,  thus,  that  heat  energy  has  two  factors — intensity 
and  quantity.  We  measure  intensity  by  means  of  a  thermometer, 
and  quantity  with  what  is  called  a  calorimeter.  The  amount  of 
heat  is  measured  by  the  rise  in  temperature  which  the  heat  pro- 
duces in  a  definite  weight  of  water.  The  unit  in  this  case  is  called 
a  calorie;  it  is  the  amount  of  heat  required  to  raise  the  temperature 
of  one  gram  of  water  one  degree  cen- 
tigrade (usually  from  15°  to  16°).  In 
technical  work  a  unit  based  on  the 
English  system  of  weights  is  used;  it 
is  called  a  British  thermal  unit  (B.t.u.) 
and  is  the  amount  of  heat  required  to 
raise  one  pound  of  water  one  degree 
Fahrenheit;  1  B.t.u.  =252  calories. 

If  it  is  desired  to  know  how  much 
heat  is  generated  when  a  certain  re- 
action takes  place,  weighed  quantities 
of  the  substances  involved  are  placed 
in  an  apparatus  surrounded  by  water, 
the  temperature  of  which  is  known. 
The  substances  are  next  allowed  to 
react,  and  as  a  result  the  heat 
generated  causes  the  temperature  of 
the  water  to  rise.  If  this  rise  is 
noted  and  the  weight  of  the  water  is 

known,  the  number  of  calories  liberated  by  the  reaction  can  be 
calculated.  The  apparatus  in  which  such  observations  are  made 
is  called  a  calorimeter.  A  convenient  form  that  is  often  used  is 
represented  by  Fig.  16;  it  is  called  a  bomb-calorimeter  because 
chemical  reactions  under  pressure  can  be  carried  out  in  it.  Sup- 
pose it  is  desired  to  know  how  much  heat  is  generated  when  char- 
coal burns.  A  weighed  amount  of  charcoal  is  put  into  the  cup, 
the  bomb  is  then  closed  and  oxygen  is  forced  in  under  pressure. 
The  reaction  is  started  by  passing  a  current  of  electricity  through 


FIG.  16. 


140  INORGANIC  CHEMISTRY  FOR  COLLEGES 

the  wire  which  passes  over  the  charcoal.  The  wire  melts  and 
ignites  the  charcoal. 

A  calorimeter  of  this  type  is  used  to  determine  the  heat  pro- 
duced when  foodstuffs  are  burned.  The  value  of  a  food  is  deter- 
mined, in  part,  by  the  amount  of  heat  which  it  produces  when  oxi- 
dized, since  one  function  of  food  is  to  furnish  heat  to  keep  up  the 
temperature  of  the  body.  One  gram  of  fat  gives  about  9000  cal- 
ories when  it  burns,  whereas  1  gram  of  starch  gives  about  4000 
calories.  Such  facts  as  these  are  of  vital  importance  in  the  science 
of  foods. 

154.  Thermochemistry. — In  all  substances  as  we  know  them, 
matter  is  associated  with  energy.  Up  to  this  point  the  changes  in 
the  matter  have  been  emphasized,  but  for  a  more  complete  under- 
standing of  these  changes  we  must  study  the  transformations  in 
energy  which  take  place  simultaneously  with  the  changes  in  matter. 
Two  substances  possessing  distinctly  different  properties  maybe 
composed  of  the  same  kind  of  matter.  A  case  which  will  be  dis- 
.  cussed  in  detail  later  can  be  mentioned  here.  A  diamond  is  a  very 
different  substance  from  a  bit  of  charcoal,  but  the  matter  in  the 
two  is  identical — they  consist  solely  of  the  element  called  carbon. 
The  great  difference  in  their  properties  can  be  traced  to  the  fact 
that  the  amount  of  energy  which  is  combined  with  the  elementary 
substance  carbon  is  different  in  the  two  cases.  This  can  be  shown 
by  a  careful  study  of  what  occurs  when  the  two  substances  are 
burned.  If  12  grams  of  charcoal  are  burned,  44  grams  of  the  gas 
carbon  dioxide  are  obtained,  the  equation  for  the  reaction  being 
C  +  (>2  =  CO2.  If  12  grams  of  diamond  are  burned,  44  grams  of 
carbon  dioxide  are  also  obtained;  the  gases  formed  in  the  two 
cases  are  identical  in  weight  and  properties.  The  study  of  the 
matter  involved  in  the  two  cases  gives  us  no  information  as  to  the 
cause  of  the  difference  between  charcoal  and  diamond.  If  the 
heat  produced  in  the  two  cases  is  determined — if  the  energy  change 
is  investigated — we  find  that  the  results  are  different.  When  1 
gram  of  charcoal  burns  the  heat  produced  is  8080  calories;  when 
1  gram  of  diamond  is  burned  7860  calories  are  set  free.  Since  the 
same  amount  of  oxygen  is  transformed  into  carbon  dioxide  in 
the  two  cases,  we  must  attribute  the  different  results  in  the  two 
cases  to  the  fact  that  carbon  as  charcoal  contains  gram  for  gram 
more  chemical  energy  than  diamond.  The  element  carbon  is  in 


THE  ENERGY  FACTOR  IN  CHEMICAL  CHANGE          141 

both  these  substances.  We  must  differentiate  between  the  element 
carbon,  and  the  elementary  substance  carbon,  which  can  appear  in  a 
variety  of  physical  forms  as  the  result  of  the  fact  that  the  element 
can  be  associated  with  different  amounts  of  chemical  energy. 
Carbon  is  present  in  carbon  dioxide,  in  wood,  in  sugar,  in  alcohol, 
and  in  thousands  of  other  compounds;  it  has  no  characteristic 
property  except  weight. 

155.  An  addition  can  be  made  to  chemical  equations  to  express 
the  amount  of  heat  set  free  when  the  reaction  indicated  by  the 
equation  takes  place.  For  example,  the  burning  of  charcoal  in 
oxygen  can  be  represented  thus: 

C  +  O2  =  CO2  +  97,000  cal. 

Such  an  expression  is  a  thermochemical  equation.  In  this  case  C 
signifies  the  number  of  grams  of  carbon  equal  to  its  atomic  weight, 
that  is,  12  grams.  Each  symbol  in  general  represents  the  weight 
in  grams  equal  to  its  atomic  weight;  this  weight  is  called  a  gram- 
atomic-weight  or  gram-atom.  In  a  thermochemical  equation 
the  formula  CC>2  stands  for  12  +  (2  X  16)  =  44  grams  of  carbon 
dioxide;  this  is  a  gram-molecular-weight  or  a  mol  of  carbon  diox- 
ide. The  above  equation  means,  then,  that  when  12  grams  of 
carbon  burn  in  oxygen,  32  grams  of  the  latter  unite  with  it  and 
form  44  grams  of  carbon  dioxide,  and,  at  the  same  time,  97,000 
calories  are  set  free;  this  number  of  calories  is  called  the  heat  of 
formation  of  carbon  dioxide. 

A  chemical  reaction  in  which  heat  is  given  off  is  said  to  be 
exothermic;  one  in  which  heat  is  absorbed  is  endothermic.  When 
an  exothermic  compound  breaks  down  into  its  elements  the 
energy  absorbed  is  equivalent  in  amount  to  that  given  off  in  its 
formation;  in  this  case  heat  is  converted  into  chemical  energy. 

Thermochemical  equations  deal  with  but  one  factor  in  the 
energy  changes  which  take  place  in  chemical  reactions — the 
quantity  factor.  For  a  complete  understanding  of  these  changes  a 
knowledge  of  the  intensity  factor  is  necessary.  We  shall  see  later 
that  the  transformation  of  chemical  energy  into  electrical  energy 
will  give  us  some  knowledge  of  this  intensity  factor. 

A  few  of  the  equations  which  have  already  been  given  will  be 
discussed  briefly  from  the  standpoint  of  energy.  The  thermo- 


142  INORGANIC  CHEMISTRY  FOR  COLLEGES 

chemical  equation  for  the  burning  of  hydrogen  in  oxygen  is  as 
follo'ws: 

2H2  +  O2  =  2H20  +  116,200  cal. 

This  means  that  when  4  grams  of  hydrogen  are  converted  into 
water-vapor,  116,200  calories  are  liberated;  1  gram  of  the  gas  pro- 
duces 29,050  calories.  The  heat  of  formation  of  carbon  dioxide 
is  97,000  calories,  this  amount  of  heat  being  produced  when  12 
grams  of  carbon  burn;  1  gram  of  carbon  furnishes,  thus,  8080 
calories.  It  is  evident  from  this  that  weight  for  weight  hydrogen 
furnishes  over  three  and  one-half  times  as  much  heat  as  carbon; 
for  this  and  other  reasons  the  gas  is  used  as  a  source  of  heat. 
Hydrogen  is  a  constituent  of  illuminating  gas,  which  is  used  as  a 
source  of  power  in  gas  engines.  The  large  amount  of  heat  pro- 
duced when  the  mixture  of  hydrogen  and  air  explodes  raises  the 
temperature  of  the  gases  to  a  high  point,  and  as  a  result  the  pres- 
sure on  the  piston  of  the  engine  is  great.  The  temperature  of 
the  hydrogen  flame  is  about  2500°,  whereas  that  of  burning  carbon 
is  about  1400°;  this  difference  is  in  part  due  to  the  fact  that  so 
much  more  heat  is  given  off  in  the  first  case,  and  as  it  is  formed 
rapidly  it  raises  the  temperature  of  the  product  of  combustion  to 
a  higher  point  than  is  attained  when  carbon  burns.  This  fact 
leads  to  some  of  the  uses  of  hydrogen  which  have  already  been 
described  (56). 

156.  It  will  be  recalled  that  the  oxides  of  certain  metals,  such  as 
mercury  and  silver,  are  decomposed  more  or  less  readily  when 
heated,  and  oxygen  is  set  free;  whereas  other  oxides  such  as  those 
of  sodium  and  calcium  can  be  heated  to  the  highest  temperatures 
attainable  without  decomposition.  A  comparison  of  the  heats  of 
formation  of  these  oxides  will  help  in  interpreting  the  facts.  The 
values  in  calories  are  as  foUows:  CaO,  145,000;  Na2O,  99,760; 
HgO,  22,000;  Ag2O,  5900.  In  each  case  the  values  are  those  of 
the  heat  set  free  when  1  atomic  weight  of  oxygen,  16  grams, 
unites  with  the  several  metals.  When  the  oxides  of  the  active 
metals  are  formed  large  amounts  of  chemical  energy  are  changed 
into  heat,  and  the  oxides  which  result  are  very  stable.  In  most 
cases  if  the  change  in  energy  when  a  compound  is  formed  is  large, 
the  compound  is  a  stable  one  toward  heat,  and  the  reverse  is  true. 

It  will  also  be  recalled  that,  in  general,  the  chloride  of  an  ele- 


THE  ENERGY  FACTOR  IN  CHEMICAL  CHANGE          143 

ment  is  not  so  readily  decomposed  by  heat  as  the  oxide;  for  exam- 
ple, silver  chloride  does  not  readily  break  down  into  silver  and 
chlorine,  whereas  silver  oxide  is  decomposed  into  its  elements  at  a 
comparatively  low  temperature.  We  can  make  oxygen  by  heating 
mercuric  oxide,  but  cannot  use  the  chloride  to  make  chlorine.  On 
the  other  hand,  platinum  and  gold,  which  are  very  inactive  metals, 
yield  chlorides  which  are  decomposed  by  heat.  The  difference  in 
stability  toward  heat  between  the  chlorides  and  oxides  can  be 
traced  to  the  fact  that  chlorine  is  a  more  active  element  than  oxy- 
gen, and  that  when  it  unites  with  other  elements,  more  chemical 
energy  is  transformed  into  heat  than  is  the  case  in  the  formation  of 
oxides.  When  1  atomic  weight  of  mercury  unites  with  oxygen 
22,000  calories  are  set  free;  when  the  same  weight  of  mercury 
unites  with  chlorine  to  form  mercuric  chloride,  HgC^,  54,490  cal- 
ories are  liberated.  In  the  case  of  Ag20  and  2AgCl  the  values 
are  5900  and  58,760,  respectively.  The  heat  of  formation  of  gold 
chloride,  AuCls,  is  22,820  calories. 

Although  the  heat  of  formation  of  a  compound  cannot  be 
taken  as  a  quantitative  measure  of  the  activity  of  the  elements 
composing  the  compound,  it  furnishes  valuable  information  in 
many  cases  as  to  the  relative  activity  of  elements  when  similar 
compounds  are  compared.  It  should  be  noted  that  certain  endo- 
thermic  compounds  are  very  stable  toward  heat.  Nitric  oxide, 
for  example,  is  formed  with  the  absorption  of  a  large  amount  of 
energy,  yet  it  resists  to  a  high  degree  decomposition  by  heat. 

The  relation  which  exists  between*  chemical  energy  and  heat  is 
considered  in  detail  in  physical  chemistry  and  cannot  be  discussed 
further  with  profit  at  this  point. 


CHAPTER  XIII 
OZONE  AND  HYDROGEN  PEROXIDE 

157.  The  fact  that  the  amount  of  energy  combined  with  a 
chemical  element  is  an  important  factor  in  determining  its  physical 
properties  has  been  brought  out  in  the  case  of  carbon,  which  can 
exist  as  charcoal  or  as  diamond.  In  the  case  of  other  elements  the 
differences  produced  in  this  way  are  even  more  marked.  Oxygen 
furnishes  an  excellent  example;  it  exists  as  oxygen  gas,  the  prop- 
erties of  which  are  familiar,  and  as  ozone,  a  substance  that  differs 
markedly  from  oxygen  in  chemical  as  well  as  physical  properties. 
Ozone  is  a  blue  gas,  has  a  marked  odor,  and  is  an  exceedingly  active 
substance;  silver  is  not  affected  by  oxygen,  but  is  blackened  by 
ozone.  Ozone  has  a  number  of  important  and  interesting  applica- 
tions, which  are  based  on  the  fact  that  it  contains  a  large  amount 
of  chemical  energy  which  can  be  readily  transformed  and,  thus, 
utilized.  The  fact  that  a  characteristic  odor  was  formed  near 
machines  which  produced  electricity  by  friction  was  observed  over 
a  hundred  years  ago.  In  1840  Schonbein  showed  that  a  distinct 
substance  was  formed,  which  could  be  made  in  other  ways.  The 
substance  was  called  ozone,  the  name  being  derived  from  the 
Greek  word  meaning  to  smell.  In  1856  Andrews  demonstrated 
the  fact  that  ozone  was  a  new  variety  of  oxygen. 

Ozone  is  formed  when  an  electric  discharge  takes  place  in  the 
air;  it  is  formed  in  this  way  during  thunder  storms  and  near  elec- 
trical machinery.  It  was  thought  at  one  time  that  ozone  is  always 
present  in  the  atmosphere,  but  recent  investigations  indicate  that 
this  conclusion  was  arrived  at  as  the  result  of  using  tests  for  the 
gas  which  were  not  characteristic.  Ozone  behaves  with  certain 
reagents  like  hydrogen  peroxide  and  certain  oxides  of  nitrogen, 
substances  which  are  present  in  the  air  in  small  proportions. 

The  formula  assigned  to  ozone  is  Os,  whereas  that  of  oxygen 
is  O2.  The  reason  for  this  will  be  given  later. 

144 


OZONE  AND  HYDROGEN  PEROXIDE 


145 


Ojone 


y  Binding  Posts 
Connecting 
with  Induction 
Coil. 
Oxygen  orAirlnfet 


f/7  Oufar 
Tube. 


on  Inner 
Tube 


158.  Preparation  of  Ozone. — The  gas  can  be  conveniently 
prepared  in  the  laboratory  by  passing  air  through  a  so-called 
ozonizer  which  is  represented  in  Fig.  17.  It  consists  of  two  con- 
centric glass  tubes  of  the  shape  indicated.  The  outside  of  the 
outer  tube  is  covered  with  tin-foil  which  is  connected  with  one 
binding-post;  the  inside  of  the  inner  tube  is  covered  in  the  same 
way  and  connected  with  the  other. 
These  binding-posts  are  joined  to  the 
terminals  of  an  induction  coil,  and 
oxygen  is  forced  through  the  space 
between  the  two  tubes.  As  the 
electrical  impulses  flow  from  the  metal 
on  one  tube  to  that  on  the  other, 
they  pass  through  the  oxygen  and  a 
part  of  it  (6  to  7  per  cent)  is  changed 
into  ozone.  Ozonizers  have  been  con- 
structed to  prepare  the  gas  on  the 
large  scale  for  commercial  use;  with 
these  air  is  used,  and  as  much  as  18 
per  cent  of  the  oxygen  contained  in 
it  is  converted  into  ozone.  The  gas  is 
also  formed  when  a  current  of  elec- 
tricity is  passed  through  dilute  sul- 
phuric acid;  the  percentage  of  the  ozone  in  the  oxygen  generated 
at  the  anode  can  be  varied  by  changing  the  strength  of  the 
current  and  the  size  of  the  electrodes. 

When  ozone  is  formed  in  the  ways  indicated  above,  electrical 
energy  is  transformed  into  chemical  energy,  which  is  taken  up  by 
the  oxygen.  As  the  result  of  increasing  the  energy  associated 
with  the  element  oxygen,  it  is  transformed  into  a  new  substance 
possessing  properties  different  from  those  of  oxygen  gas.  The 
energy  required  to  change  oxygen  to  ozone  has  been  determined 
experimentally.  When  this  is  expressed  as  heat  energy  we  can 
represent  it  by  the  following  equation: 

3O2  +  68,200  cal.  <=»  2O3 

When  ozone  decomposes  into  oxygen  this  energy  is  set  free;  as  a 
consequence,  ozone  is,  as  we  might  expect,  a  more  active  oxidizing 
agent  than  free  oxygen. 


FIG.  17. 


146  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Ozone  is  formed  in  small  quantities  when  potassium  chlorate 
is  heated.  It  will  be  recalled  that  oxygen  is  prepared  in  this  way. 
From  the  standpoint  of  atoms  it  is  possible  that  the  changes 
involved  are  represented  by  the  following  equations: 

KC1O3  =  KC1  +  3O 
2O  =  O2 
3O  =  O3 

Oxygen  atoms  are  first  liberated  and  then  unite  to  form  oxygen 
gas  and  ozone.  At  the  temperature  at  which  the  reaction  occurs, 
oxygen  is  much  more  stable,  and  consequently  is  formed  in  the 
larger  quantity.  That  ozone  is  present  in  the  oxygen  generated 
by  heating  potassium  chlorate  can  be  shown  by  exposing  to  it  a 
piece  of  paper  on  which  has  been  put  a  solution  of  starch  containing 
potassium  iodide;  ozone  liberates  iodine,  which  produces  a  blue 
color  with  starch. 

Ozone  is  also  formed  when  certain  elements,  such  as  phos- 
phorus and  zinc,  oxidize  slowly  in  moist  air.  It  is  probable  that 
molecules  of  oxygen,  O2,  unite  directly  with  these  elements  and 
that  the  resulting  oxides  are  unstable  and  break  down,  giving  the 
ordinary  oxides  and  oxygen  atoms,  some  of  which  unite  to  form 
ozone. 

159.  Physical  Properties  of  Ozone. — Ozone  is  a  pale-blue  gas 
which  has  a  characteristic  odor.     It  can  be  obtained  as  a  blue 
liquid,  which  boils  at  —119°,  by  passing  a  mixture  of  ozone  and 
oxygen  through  a  vessel  surrounded  by  liquid  oxygen.     It  is  much 
more  soluble  in  water  than  is  oxygen;    100  volumes  of  water  dis- 
solve at  0°,  50  volumes  of  ozone  and  but  4  volumes  of  oxygen.    One 
liter  of  ozone  at  0°  and  760  mm.  pressure  weighs  2.144  grams. 

160.  Chemical  Conduct  of  Ozone. — Ozone  decomposes  slowly 
at  ordinary  temperatures  into  oxygen;  in  the  presence  of  finely 
divided  platinum  and  at  about  250°  the  change  is  very  rapid.     It 
is  unsafe  to  keep  liquid  ozone  because  it  may  decompose  with 
explosive  violence. 

Ozone  is  absorbed  by  turpentine  and  other  oils  and  forms  com- 
pounds with  them;  it  changes  rubber  into  a  substance  which  is 
inelastic. 


OZONE  AND  HYDROGEN  PEROXIDE  147 

The  activity  of  ozone  has  been  noted;  it  converts  silver  into 
silver  peroxide,  which  is  black  in  color: 

2O3  +  2Ag  =  Ag2O2  +  2O2 

The  reaction  can  be  used  as  a  test  for  ozone  provided  no  sulphur 
compound  which  blackens  silver  is  present. 

Ozone  liberates  iodine  from  potassium  iodide: 

O3  +  2KI  +  H2O  =  2KOH  +  I2  +  O2 

A  solution  containing  potassium  iodide  and  starch,  which  gives  a 
blue  color  with  free  iodine,  is  used  in  showing  the  presence  of 
ozone,  but  as  many  active  oxidizing  agents  act  in  this  way  the 
reaction  cannot  be  used  as  a  test  for  the  gas.  From  an  examina- 
tion of  the  last  two  equations  it  will  be  seen  that  when  ozone  oxi- 
dizes substances,  but  one  atom  of  the  oxygen  in  each  molecule  of 
the  gas  unites  with  the  substance  oxidized,  and  that  oxygen  gas  is 
formed  from  the  remaining  two  atoms.  This  usually  occurs  when 
ozone  acts  as  an  oxidizing  agent. 

The  manner  in  which  the  atoms  are  linked  together  in  oxygen 
and  in  ozone  is  an  interesting  subject  for  speculation.  The 
graphic  formulas  usually  assigned  to  these  substances  are  as 
follows : 


0=0 


These  formulas  indicate  that  each  oxygen  atom  has  the  valence 
2.  Another  formula  for  ozone,  which  expresses,  perhaps,  its  prop- 
erties more  adequately  is  this:  O=O=O.  We  can  understand 
how  a  molecule  made  up  in  this  way  could  break  down  into  a 
molecule  of  oxygen  O2,  and  an  oxygen  atom  which  would  oxidize 
other  substances;  such  a  decomposition  would  be  represented  thus: 

O=O=O  =  O=O  +  0 

The  formula  indicates  that  one  of  the  oxygen  atoms  has  the 
valence  4,  a  view  for  which  there  is  much  independent  evidence. 

161.  Allotropic  Modifications  of  Elements. — Oxygen  gas  and 
ozone  are  said  to  be  allotropic  modifications  of  the  element  oxygen; 


148  INORGANIC  CHEMISTRY  FOR  COLLEGES 

the  term  is  used  to  express  the  fact  that  the  same  chemical  element 
exists  as  two  distinct  substances;  they  have  different  physical 
and  chemical  properties,  and  equal  weights  of  the  two  possess 
different  amounts  of  chemical  energy.  These  differences  are  the 
result  of  the  fact  that  they  possess  different  atomic  structures. 
Many  elements  exist  in  two  or  more  allo tropic  modifications;  the 
case  of  carbon  in  the  form  of  diamond  and  charcoal  has  already 
been  noted.  Gaseous  oxygen  and  liquid  oxygen  are  not  allotropic 
forms;  the  meaning  of  the  word  is  limited  to  those  cases  where  the 
two  substances  exist  in  the  same  physical  state;  the  two  forms 
must  be  both  gases,  both  liquids,  or  both  solids. 

162.  Uses  of  Ozone. — The  fact  that  ozone  is  an  active  oxidizing 
agent  has  led  to  many  applications  of  the  gas.  It  will  be  recalled 
that  bleaching,  in  many  cases,  is  the  result  of  oxidation.  When 
hypochlorous  acid,  HC1O,  is  used  for  this  purpose,  hydrochloric 
acid  is  left  behind  in  the  fabric  bleached.  This  prevents  in  certain 
cases  the  use  of  this  cheap  and  efficient  agent.  When  ozone 
bleaches,  the  by-product  is  oxygen,  a  gas  which  escapes  and  is  not 
deleterious;  for  this  reason  it  is  of  particular  value  in  bleaching 
substances  which  are  affected  by  active  chemical  reagents.  It  has 
been  found  of  great  value  in  bleaching  flour,  starch,  oils,  waxes, 
ivory,  wool,  and  silk.  Like  other  active  oxidizing  agents,  ozone 
destroys  bacteria.  Water  is  sterilized  for  household  use  by  allowing 
it  to  come  into  contact  with  air  that  has  been  passed  through  an 
ozonizer.  This  method  is  used  to  some  extent  in  purifying  the 
water  supplies  of  several  European  cities;  in  America,  bleaching 
powder  or  chlorine  have  found  a  limited  use  for  this  purpose,  being 
preferred  to  ozone  on  account  of  their  lower  cost. 

Ozone  is  also  used  to  purify  and  deodorize  the  air  of  rooms  in 
which  large  numbers  of  people  congregate.  A  number  of  theaters 
in  Germany  are  equipped  with  apparatus  to  introduce  ozonized  air 
into  the  ventilating  system.  It  has  not  yet  been  definitely  estab- 
lished whether  ozone  destroys  the  bacteria  and  substances  pos- 
sessing disagreeable  odors  that  are  found  in  the  air  of  a  room  filled 
with  people;  it  is  possible  that  the  ozone  merely  affects  the  sense 
of  smell  to  such  an  extent  that  the  presence  of  the  unpleasant 
odors  is  not  observed.  The  fact  has  been  apparently  established 
that  ozone  acts  as  an  oxidizing  agent  only  in  the  presence  of  appre- 
ciable amounts  of  water-vapor. 


OZONE  AND  HYDROGEN  PEROXIDE  149 

Ozone  is  said  to  have  a  marked  effect  on  tobacco  smoke  sus- 
pended in  the  air,  and  its  use  in  ventilation  for  clearing  away  the 
substances  in  smoke  which  affect  the  eyes  and  throat  has  been 
suggested. 

HYDROGEN  PEROXIDE 

163.  In  1818  Thenard,  when  studying  the  action  of  acids  on  the 
oxides  of  metals,  discovered  that  certain  of  these  yielded  a  com- 
pound of  hydrogen  and  oxygen  which  possessed  remarkable  prop- 
erties; the  substance  was  shown  to  be  a  liquid  and  to  have  the 
composition  represented  by  the  formula  H2O2;  it  is  now  called 
hydrogen  peroxide.  Most  oxides  of  metals  react  with  acids  to 
form  water  and  salts;  the  reaction  between  barium  oxide  and 
hydrochloric  acid  is  typical: 

BaO  +  2HC1  -  H2O  +  BaCl2 

When  barium  peroxide  is  used,  however,  hydrogen  peroxide,  and 
not  water,  is  formed : 

BaO2  +  2HC1  =  H2O2  +  BaCl2 

A  study  of  the  dioxides  of  the  metals  has  led  to  the  conclusion  that 
they  can  be  divided  into  two  classes:  those  that  yield  hydrogen 
peroxide  when  treated  with  acids,  and  those  that  do  not.  Exam- 
ples of  the  latter  class  have  already  been  met  with.  When  man- 
ganese dioxide  reacts  with  hydrochloric  acid,  the  products  are 
water  and  manganese  tetrachloride,  which  breaks  down  into 
chlorine  and  manganese  chloride: 

MnO2  +  4HC1  =  MnCl4  +  2H2O 
MnCU  =  MnCl2  +  C12 

Manganese  in  MnO2  has  the  valence  4  and  is  first  converted  into  a 
chloride  in  which  the  valence  of  the  metal  is  4;  at  room  tem- 
perature 2  atoms  of  chlorine  break  off  from  the  compound,  and  in 


150  INORGANIC  CHEMISTRY  FOR  COLLEGES 

the  resulting  compound  the  metal  has  the  valence  2.     The  graphic 
formulas  of  these  compounds  are  represented  as  follows: 


y> 

Mn/ 
N) 


cl 

Mn 


All  metallic  oxides  which  contain  2  oxygen  atoms  and  behave  in 
this  way  are  called  dioxides.  Oxides  which  yield  hydrogen  peroxide 
when  treated  with  an  acid  are  called  peroxides.  The  formula  of 
barium  peroxide,  BaCb,  leads  to  the  conclusion  that  in  it  the 
valence  of  barium  is  4;  but  this  is  not  the  case.  The  difference  in 
the  behavior  of  barium  peroxide  and  manganese  dioxide  is  explained 
on  the  assumption  that  in  the  former  compound  the  valence  of 
barium  is  2.  The  oxygen  atoms  are  joined  to  the  barium  as  indi- 
cated by  one  of  the  following  formulas: 


or       a== 


If  this  view  is  correct,  the  formula  of  hydrogen  peroxide  is 

H— 0  H 

or 
H— O  H 

The  formulas  of  the  peroxides  involving  an  oxygen  atom  with  the 
valence  4  represent  better  the  chemical  behavior  of  these  sub- 
stances. 

164.  Occurrence  of  Hydrogen  Peroxide. — Minute  quantities 
of  hydrogen  peroxide  have  been  reported  as  being  present  in  rain 
and  snow.  It  is  possible  that  the  substance  is  formed  from  water 
and  the  oxygen  in  the  air  in  the  presence  of  sunlight;  the  bleaching 
of  moist  linen  when  exposed  to  the  sunlight  has  been  explained  on 
the  hypothesis  that  the  hydrogen  peroxide  formed  in  this  way 
oxidizes  the  coloring  matters  in  the  unbleached  cloth.  Enough 
hydrogen  peroxide  to  be  detected  is  formed  when  certain  metals,  such 
as  zinc  and  lead,  oxidize  in  moist  air.  It  is  probable  that  when 
substances  are  oxidized  by  gaseous  oxygen,  molecules  of  the  latter 


OZONE  AND  HYDROGEN  PEROXIDE  151 

first  add  directly  to  them,  and  peroxides,  which  are  more  or  less 
stable,  are  formed.  In  the  case  of  sodium,  the  peroxide  Na2O2  is 
stable  and  is  formed  when  the  substance  is  burned  in  oxygen.  It 
is  probable  that  the  peroxides  of  most  of  the  metals,  being  unstable, 
break  down  and  yield  the  ordinary  oxide  and  traces  of  ozone  or 
hydrogen  peroxide. 

165.  Preparation  of  Hydrogen  Peroxide. — Hydrogen  peroxide 
is  prepared  by  stirring  finely  powdered  barium  peroxide  into  a 
dilute  solution  of  sulphuric  acid,  which  is  kept  cold  with  ice: 

BaO2  +  H2SO4  =  H202  +  BaS04 

The  product  is  filtered  and  diluted  with  enough  water  to  make  it  a 
3.5  per  cent  solution.  Hydrogen  peroxide  decomposes  more  or  less 
rapidly  in  the  presence  of  bases  and  salts,  and  for  this  reason  a 
trace  of  acid  is  often  left  in  the  solution.  Certain  organic  sub- 
stances are  thought  to  act  as  so-called  negative  catalyzers,  that  is, 
they  retard  decomposition;  acetanilide  is  used  for  this  purpose  in 
certain  commercial  brands  of  hydrogen  peroxide.  Sulphuric  acid 
is  sometimes  replaced  by  hydrochloric  acid  or  phosphoric  acid  in 
preparing  hydrogen  peroxide;  the  latter  yields  a  solution  which  is 
supposed  to  keep  better  than  that  obtained  when  sulphuric  acid  is 
used. 

Hydrogen  peroxide  is  also  made  commercially  from  ammonium 
persulphate  which  is  prepared  by  the  electrolysis  of  ammonium 
sulphate  (311).  Hydrogen  peroxide  is  formed  when  hydrogen 
burns  in  air  or  oxygen,  although  at  the  high  temperature  at  which 
the  reaction  takes  place  it  decomposes  completely  unless  special 
precautions  are  taken.  If  a  jet  of  burning  hydrogen  is  held  in 
contact  with  ice  for  a  short  time  and  the  water  formed  is  examined, 
it  will  be  found  to  contain  enough  hydrogen  peroxide  to  give  a  dis- 
tinct test.  It  is  probable  that  a  molecule  of  hydrogen  and  one 
of  oxygen  unite  directly  to  form  hydrogen  peroxide: 

H2  +  02  =  H202 

It  is  probable,  as  has  been  stated,  that  when  oxygen  reacts  with 
another  substance  the  first  action  consists  in  direct  combination  of 
the  two;  in  most  cases,  however,  the  peroxides  formed  are  un- 
stable and  decompose — a  fact  which  is  illustrated  in  the  case  of 
the  union  of  hydrogen  and  oxygen. 


152  INORGANIC  CHEMISTRY  FOR  COLLEGES 

166.  Physical  Properties  of  Hydrogen  Peroxide. — Hydrogen 
peroxide  is  a  syrupy  liquid,  which  has  the  density  1.46.     It  boils  at 
69°  when  the  pressure  is  reduced  to  26  mm.,  and  can  be  frozen  to 
a  crystalline  solid  which  melts  at  —2°.     The  substance  is  seldom 
isolated  on  account  of  the  fact  that  it  is  apt  to  decompose  with 
explosive  violence.     It  is  usually  kept  in  solution  in  water.     The 
solution  ordinarily  sold  contains  3  per  cent  of  hydrogen  peroxide; 
its  strength  is  often  expressed  in  "  volumes  ";  a  10- volume  solution 
will  yield  ten  times  its  volume  of  oxygen  when  decomposed.     By 
evaporating  a  3  per  cent  solution  of  hydrogen  peroxide  at  70° 
the  strength  can  be  increased  to  45  per  cent  without  much  loss. 

167.  Chemical  Conduct  of  Hydrogen  Peroxide. — When  hydro- 
gen peroxide  decomposes  into  water  and  oxygen  a  large  amount  of 
energy  is  set  free : 

2H2O2  =  2H2O  +  O2  +  46,200  cal. 

This  fact  leads  to  the  conclusion  that  hydrogen  peroxide  is  an 
active  oxidizing  agent.  The  decomposition  takes  place  slowly  at 
ordinary  temperatures,  the  rate  being  markedly  affected  by  rise  in 
temperature  and  the  presence  of  catalyzers  such  as  charcoal,  man- 
ganese dioxide,  and  finely  divided  platinum,  silver  or  gold.  Hydro- 
gen peroxide  liberates  iodine  from  potassium  iodide;  the  reaction 
recalls  the  behavior  of  ozone  with  this  substance : 

H2O2  +  2KI  =  2KOH  +  I2 

A  solution  containing  potassium  iodide  and  starch  is  used  for  test- 
ing for  hydrogen  peroxide;  the  formation  of  a  blue  color  is  not 
definite  proof  of  the  presence  of  hydrogen  peroxide,  however,  since 
ozone  produces  the  same  effect.  Hydrogen  peroxide  does  not 
blacken  silver,  but  ozone  does. 

Hydrogen  peroxide  acts  with  certain  bases  as  an  acid;  when  a 
solution  of  barium  hydroxide  is  treated  with  it  barium  peroxide 
and  water  are  formed : 

Ba(OH)2  -f  H2O2  =  BaO2  +  2H2O 

The  precipitate  formed  is  a  hydrate  of  the  composition  BaO2,8H2O. 
The  hydroxides  of  certain  other  elements  act  in  the  same  way. 
Some  of  the  hydrated  peroxides  made  in  this  way  are  used  in 


OZONE  AND  HYDROGEN  PEROXIDE  153 

pharmacy;  zinc  peroxide,  which  is  a  constituent  of  salves,  is 
believed  to  have  antiseptic  properties. 

It  is  a  remarkable  fact  that  with  certain  substances  hydrogen 
peroxide  acts  as  a  reducing  agent;  for  example,  silver  oxide  is 
reduced  to  metallic  silver: 

Ag2O  +  H2O2  =  2Ag  +  H2O  +  O2 

It  is  probable  that  a  peroxide  is  first  formed  which,  being  unstable, 
breaks  down  spontaneously  into  silver  and  oxygen.  Certain  other 
substances  which  contain  a  high  percentage  of  oxygen  behave  in 
the  same  way.  When  a  solution  of  potassium  dichromate  is  added 
to  one  of  hydrogen  peroxide  containing  a  little  sulphuric  acid,  a 
blue  color  is  developed,  which  soon  disappears.  If  the  solution  is 
immediately  shaken  with  a  small  amount  of  ether,  the  blue  sub- 
stance dissolves  in  the  latter  and  persists  for  a  longer  time.  This 
behavior  of  hydrogen  peroxide  is  used  as  a  test  for  it.  It  is  prob- 
able that  the  unstable  blue  substance  formed  is  a  perchromic  acid 
which  contains  a  high  percentage  of  oxygen. 

168.  Uses  of  Hydrogen  Peroxide. — The  uses  to  which  hydrogen 
peroxide  is  put  are  based  upon  the  fact  that  it  is  an  oxidizing  agent; 
since  only  water  is  formed  from  it  as  a  by-product,  it  is  of  particular 
value  in  bleaching  substances  which  cannot  be  treated  with  hypo- 
chlorites.  It  is  used  as  a  bleaching  agent  for  materials  of  an  animal 
source,  such  as  wool,  hair,  feathers,  silk,  bone,  and  ivory.  It  was 
formerly  much  used  as  a  hair  bleach  for  toilet  use.  In  bleaching 
with  hydrogen  peroxide  the  latter  is  not  often  isolated.  A  solution 
is  prepared  by  adding  sodium  peroxide  to  a  dilute  solution  of  an 
acid  which  is  kept  cold: 

Na2O2  +  2HC1  =  2NaCl  +  H2O2 

Hydrogen  peroxide  is  much  used  in  the  household  as  an  anti- 
septic and  disinfectant  on  account  of  the  fact  that  oxygen  is  liber- 
ated when  it  is  brought  into  contact  with  blood  and  decaying 
organic  matter;  it  is  doubtful,  however,  whether  it  is  of  much  value 
for  this  purpose. 


154  INORGANIC  CHEMISTRY  FOR  COLLEGES 


EXERCISES 

1.  What  volume  of  ozone  could  be  obtained  from  100  c.c.  of  oxygen  if 
the  conversion  were  complete? 

2.  Ten  liters  of  air  were  passed  through  an  ozonizer  and  10  per  cent  of 
the  oxygen  was  converted  into  ozone.     What  was  the  total  volume  of  the 
gases  obtained? 

3.  Calculate  approximately  the  volume  of  oxygen  which  is  set  free  when 
100  grams  of  a  3  per  cent  solution  of  hydrogen  peroxide  decomposes  into 
water  and  oxygen.     How  could  you  express  the  "strength"  of  the  solution 
in  volumes? 

4.  State  two  ways  of  obtaining  from  a  mixture  of  oxygen  and  ozone  con- 
taining 10  per  cent  of  the  latter,  a  mixture  of  the  two  gases  which  contained 
a  higher  percentage  of  ozone. 

5.  Calculate  from  the  formula  of  ozone  the  weight  of  1  liter  of  the  gas  at 
0°  and  760  mm. 

6.  Calculate  the  percentage  of  ozone  in  a  mixture  of  oxygen  and  ozone 
1  liter  of  which  weighs  1.572  grams  at  0°  and  760  mm. 


CHAPTER  XIV 

PROPERTIES  OF  GASES,  LIQUIDS,  AND  SOLIDS 

169.  We  have  seen  that  a  study  of  the  changes  in  energy  which 
take  place  in  chemical  reactions  adds  materially  to  our  knowledge 
of  matter.     Energy  produces  also  what  we  have  called  physical 
changes.     These  are  of  great  importance  and  must  be  treated  in 
some  detail,  although  they  are  rightly  a  part  of  physics.     Physics 
and  chemistry  are  so  closely  interwoven  that  the  fundamentals 
of  one  science  must  be  understood  to  appreciate  the  other;    we 
make  use  of  physical  changes  in  matter  to  recognize  and  interpret 
chemical  phenomena. 

The  physical  effect  of  heat  energy  on  matter  is  striking;  we 
know  that  we  can  change  water  into  steam — a  gas — and  into  ice — 
a  solid.  Such  changes  as  these  are  used  in  studying  matter  from 
the  standpoint  of  chemistry,  and  in  formulating  theories  as  to  the 
composition  of  molecules.  We  must,  accordingly,  have  a  definite 
knowledge  of  the  physical  properties  of  matter  if  we  are  to  use 
them  in  helping  to  solve  our  chemical  problems. 

170.  Gases. — The  characteristic  property  of  a  gas  is  that  of 
diffusion;    a  gas  distributes  itself  uniformly  in  any  space  into 
which  it  is  put;   its  volume  is  determined  by  the  volume  of  the 
vessel  which  contains  it.     We  make  use  of  the  fact  that  a  charac- 
teristic property  of  a  gas  is  diffusion  when  we  use  ozone  to  purify 
the  air  in  a  room.     It  is  only  necessary  to  liberate  some  of  the  gas 
at  one  point,  because  it  diffuses  and  finally  becomes  uniformly 
distributed. 

Gases  are  very  compressible,  that  is,  a  relatively  small  change 
in  pressure  on  a  gas  has  a  marked  effect  on  its  volume.  The  extent 
to  which  they  expand  when  heated  is  great  compared  with  that  of 
liquids  and  solids.  These  facts  have  led  to  a  conception  of  the 
structure  of  gases  which  has  been  valuable  in  science.  According 
to  this  theory  gases  are  made  up  of  small  particles,  called  mole- 

155 


156  INORGANIC  CHEMISTRY  FOR  COLLEGES 

cules,  which  are  in  constant  motion.  The  actual  space  occupied 
by  the  molecules  is  small  as  compared  with  the  space  occupied  by 
the  gas.  This  is  evident  from  the  fact  that  the  volume  of  a  sub- 
tance  in  the  form  of  a  gas  is  much  greater  than  its  volume  when 
in  the  form  of  a  liquid.  When  water  is  changed  into  steam,  at 
atmospheric  pressure,  the  volume  of  the  latter  is  over  1500  times 
the  volume  of  the  water  from  which  the  steam  was  produced. 

The  particles  in  a  gas  move  about  freely  and  collide  with  one 
another  and  the  sides  of  the  vessel  which  contains  them.  These 
collisions  produce  the  pressure  which  the  gas  exerts.  There  is  no 
loss  in  energy  in  the  gas  as  the  result  of  these  collisions.  When 
two  pieces  of  matter  strike  each  other  a  part  of  the  energy  of 
motion  is  converted  into  heat.  If  this  took  place  in  a  gas  it  would 
slowly  lose  energy,  for  the  heat  produced  would  be  taken  up  by 
the  surroundings,  and  the  temperature  of  the  gas  would  fall. 
Since  this  does  not  occur,  the  molecules  are  considered  to  be 
perfectly  elastic,  that  is,  when  they  collide  no  energy  of  motion  is 
changed  into  heat;  the  particle  starts  off  after  the  collision  with 
the  same  energy  it  had  before.  This  conception  of  the  physical 
structure  of  a  gas  is  known  as  the  kinetic  theory  of  gases,  on 
account  of  the  fact  that  it  postulates  moving  particles. 

The  study  of  the  physical  properties  of  gases  has  led  to  results 
that  make  it  possible  to  determine  the  size  of  molecules.  One 
gram-molecular-weight  of  hydrogen  (2  grams  which  occupy  22.4 
liters)  contains  under  standard  conditions  6.16  X  1023  molecules.1 
Each  molecule  of  hydrogen  weighs  0.02388  grams. 

171.  Liquids. — Liquids  are  characterized  by  possessing  a 
definite  volume  but  no  fixed  form;  they  take  the  shape  of  the  vessel 
containing  them,  except  when  in  minute  quantities,  as  drops,  they 
assume  a  form  more  or  less  spherical.  In  many  cases  one  liquid 
will  mix  completely  with  another;  in  such  mixtures  the  constituents 
are  evenly  distributed.  Although  one  liquid  may  be  much  heavier 
than  the  other,  the  mixture  does  not  separate  into  layers.  In 
other  cases  liquids  are  immiscible,  such  as  oil  and  water.  No 
adequate  theory  of  the  structure  of  liquids  has  been  put  forward. 
According  to  the  kinetic  theory  the  molecules  of  a  liquid  are 
packed  close  together  and  attract  one  another;  the  attraction  is 

1  10"  is  a  brief  way  of  expressing  the  number  made  up  of  1  followed  by 
23  ciphers. 


PROPERTIES  OF  GASES,  LIQUIDS,  AND  SOLIDS  157 

not  great  enough,  however,  to  prevent  their  motion.  There  is  not 
much  free  space  between  the  molecules  because  liquids  are  very 
incompressible — a  great  pressure  has  little  effect  in  decreasing  the 
volume  of  a  liquid.  If  the  pressure  on  a  gas  is  doubled,  its  volume 
decreases  one-half;  the  change  in  volume  of  water — a  liquid — 
when  the  pressure  is  changed  from  1  to  2  atmospheres  is  0.05  per 
cent. 

172.  Solids. — Solids  possess  a  definite  form.  According  to  the 
kinetic  theory  the  molecules  are  closely  packed  and  the  attraction 
between  them  is  great  enough  to  prevent  free  motion.  It  is  pos- 
sible in  a  solid,  therefore,  to  have  a  definite  arrangement  of  the 
molecules  which  remains  fixed;  this  is  seen  in  the  fact  that  many 
solids  assume  definite  forms  which  are  characteristic.  For  example, 
when  sodium  chloride  separates  from  a  solution  in  water  it  appears 
as  cubes;  the  molecules  in  building  up  the  solid  arrange  them- 
selves in  a  definite  mathematical  way  and  a  crystal  is  formed. 
Other  substances  appear  in  other  geometrical  forms;  alum,  for 
example,  crystallizes  in  regular  octahedra,  which  have  eight 
triangular  faces.  A  great  variety  of  solid  forms  can  be  built  up  by 
combining  a  number  of  planes  which  cut  each  other  at  different 
angles.  For  each  substance  the  angles  at  which  the  surfaces  meet 
are  characteristic.  The  study  of  the  forms  of  crystals  has  devel- 
oped into  a  science  called  crystallography,  a  knowledge  of  which  is 
of  service  to  the  chemist. 

173.  Crystallography. — The  great  number  of  different  forms  of  crystals 
that  exist  are  classified  by  considering  them  as  built  up  upon  axes.  The 
crystals  in  the  regular  system  are  referred  to  three  axes  of  equal  length  which 
cut  one  another  at  right  angles.  If  planes  are  passed  through  the  ends  of 
the  axes  and  perpendicular  to  them,  a  cube  is  formed;  if  they  are  passed  so 
that  each  plane  cuts  three  axes,  the  solid  formed  has  eight  faces  and  is  called 
an  octahedron.  In  the  tetragonal  system  the  axes  are  at  right  angles  and  two" 
of  them  are  equal  in  length.  When  planes  are  passed  through  the  axes  in 
the  way  described  above,  a  parallelepiped  is  formed  in  the  first  case  and,  in 
the  second,  a  figure  composed  of  two  square  pyramids  joined  at  their  bases. 
The  crystals  in  the  rhombic  system  are  built  up  on  three  axes  of  unequal 
length,  which  cut  one  another  at  right  angles.  In  the  monoclinic  system  there 
are  two  axes  at  right  angles  and  a  third  which  intersects  one  of  these  at  right 
angles  and  is  inclined  toward  the  other;  the  axes  may  vary  in  length.  Crys- 
tals in  the  triclinic  system  are  built  up  on  three  axes  which  may  differ  in  length 
and  in  the  angles  at  which  they  cut  one  another.  In  the  hexagonal  system 
there  are  three  axes  in  the  same  plane  which  cut  one  another  at  an  angle 


158  INORGANIC  CHEMISTRY  FOR  COLLEGES 

of  60°,  and  a  fourth  axis  which  passes  through  their  intersection  and  is  per- 
pendicular to  the  plane  in  which  the  other  axes  lie.  This  system  yields 
hexagonal  prisms  and  pyramids. 

It  is  readily  seen  that  a  great  variety  of  crystal  forms  can  be  built  up 
in  the  way  indicated  above.  Crystals  often  show  the  faces  of  two  or  more 
forms;  thus  in  the  regular  system  the  crystal  may  occur  as  an  octahedron, 
the  six  corners  of  which  have  been  removed  by  passing  planes  through  the 
crystal  perpendicular  to  the  axes;  such  a  form  is  a  combination  of  an  octa- 
hedron and  a  cube,  and  the  crystal  has,  accordingly,  fourteen  faces — eight 
furnished  by  the  octahedral  form  and  six  by  the  cube.  Great  complexity 
can  arise  in  this  way  and  the  subject  can  be  followed  only  with  the  aid  of 
models. 

Many  solids  do  not  apparently  possess  the  regular  arrangement  of  mole- 
cules which  is  thought  to  be  present  in  substances  that  form  crystals;  such 
solids  are  said  to  be  amorphous,  the  word  being  derived  from  the  Greek  word 
signifying  without  form.  Most  crystalline  substances  have  a  definite  melting- 
point,  that  is,  they  change  from  a  solid  to  a  liquid  at  a  definite  temperature; 
amorphous  substances  either  decompose  before  they  melt,  or  become  slowly 
viscous  as  the  temperature  rises — there  is  no  temperature  just  above  which 
the  substance  is  a  liquid,  and  just  below  which  it  is  a  solid. 

174.  Properties   Common  to   Gases,  Liquids,  and   Solids. — 

The  characteristic  physical  properties  of  the  three  states  of  matter, 
we  have  seen,  are  diffusibility  for  gases,  indeterminate  form  for 
liquids,  and  definite  form  for  solids.  Seeking  an  explanation  for 
this  we  come  to  the  conclusion  that  whether  a  substance  is  a  solid, 
liquid,  or  gas  is  determined  by  the  mobility  of  its  molecules.  All 
matter  is  made  up  of  molecules  and  has,  accordingly,  properties 
which  are  common  to  the  three  states  of  matter.  Many  such 
properties  have  been  studied  with  care.  Only  those  properties 
which  are  used  by  the  chemist  in  identifying  substances  or  in  inter- 
preting chemical  transformations  will  be  considered  at  length  here. 
Many  properties  are  carefully  studied  in  engineering  and  other 
applications  of  physics.  For  example,  what  is  known  as  tensile 
strength  is  one  of  the  properties  which  determine  what  metal  should 
be  used  in  building  a  bridge;  the  electrical  conductivity  of  copper 
is  the  property  which  leads  to  its  selection  as  a  material  from  which 
to  make  telephone  wires.  Such  properties  are  discussed  in  detail 
in  physics.  The  chemist  uses  density  as  we  shall  see,  in  arriving  at 
a  system  of  atomic  weights,  and  the  determination  of  the  boiling- 
point  of  a  liquid  is  the  simplest  practical  way  to  identify  it.  These, 
and  other  properties  of  like  significance,  must  be  familiar  to  the 
student  of  chemistry  and  will  now  be  discussed. 


PROPERTIES  OF  GASES,  LIQUIDS,  AND  SOLIDS          159 

175.  Density. — The  weight  in  grams  of  1  c.c.  of  a  substance  is 
called  its  density.     On  account  of  the  fact  that  the  density  of 
gases  is  so  small,  the  weight  of  1  liter  of  the  gas  at  0°  and  760  mm. 
pressure  is  often  used  for  this  quantity.     The  density  of  water  at 
4°  is  1,  of  iron  7.84,  of  aluminium  2.6,  of  sodium  chloride  2.13,  etc. 
The  numbers  are  of  great  value  for  from  them  we  learn  the  relative 
heaviness   of  substances.     Since  the  volumes   of  all  substances 
change  with  change  in  temperature  and  pressure,  and  since  density 
is  the  weight  of  a  unit  volume,  the  density  of  a  substance  is  dif- 
ferent under  different  conditions.     In  recording  density,  there- 
fore, the  temperature  and  pressure  should  always  be  given  unless 
the  value  is  that  of  the  substance  under  the  conditions  accepted  as 
standard,  which  are  0°  and  760  mm.     The  change  in  the  density 
of  a  gas  under  varying  conditions  of  temperature  and  pressure  is 
of  such  importance  that  it  has  already  been  discussed  at  great 
length  in  Chapter  VIII. 

176.  The  term  specific  gravity  is  much  used.     The   specific 
gravity  of  a  substance  is  a  number  which  expresses  the  relation 
between  the  weight  of  a  substance  and  the  weight  of  an  equal 
volume  of  some  other  substance.     For  example,  if  the  liquid  which 
filled  a  certain  flask  weighed  20  grams,  and  the  water  which  filled 
the  same  flask  weighed  10  grams,  then  the  specific  gravity  of  the 
liquid   compared   with   water   is   20  -r-  10  =  2.     The   expression 
specific  gravity  is  a  bad  one,  for  its  value  is  not  a  specific,  but  a 
relative  number.     The  specific  gravity  of  solids  and  liquids  is 
referred  to  water  as  the  standard;   for  gases,  hydrogen  or  air  is 
used.    When  water  at  4°  is  the  standard,  the  specific  gravity  is 
equal  to  the  density,  since  1  c.c.  of  water  at  4°  weighs  1  gram. 

177.  Specific  Heat. — When  heat  energy  is  supplied  to  matter  it 
is  absorbed,  and  as  a  consequence  the  temperature  of  the  body 
affected  rises.     The  specific  heat  of  a  substance  is  the  number  of 
calories  required  to  raise  the  temperature  of  1  gram  of  it  1  degree 
centigrade.     A  calorie  is  defined  as  the  amount  of  heat  required  to 
raise  1  gram  of  water  1  degree  (15°-16°);    the  specific  heat  of 
water  at  15°  is  1,  accordingly.     The  specific  heats  of  substances 
vary  greatly;  they  are  all  less  than  1.    A  few  values  are  as  follows: 
iron,  0.112;    zinc,  0.093;    gold,  0.032;    yellow  phosphorus,  0.19. 
The  specific  heat  varies  with  temperature  and  pressure,  and  in 
recording  the  value  for  a  substance  these  conditions  must  be  stated. 


160  INORGANIC  CHEMISTRY  FOR  COLLEGES 

178.  The  Boiling-point. — If  water  is  left  undisturbed  in  an  open 
vessel  it  will  slowly  disappear,  evaporate;   this  is  due  to  the  fact 
that  the  liquid  changes  into  water-vapor,  a  gas,  which  mixes  with 
the  air.     The  word  vapor  is  commonly  applied  to  the  gaseous 
form  of  a  substance  which  at  ordinary  temperatures  is  a  liquid; 
vapors  are  gases  and  their  behavior  is,  in  general,  in  accord  with 
the  gas  laws. 

We  have  seen  that  the  gas  given  off  from  water  exerts  a  pressure 
(89)  and  that  this  increases  with  rise  in  temperature.  When 
the  pressure  of  the  vapor  from  any  liquid  just  exceeds  that  of  the 
air,  the  liquid  boils;  at  760  mm.  pressure  water  boils  at  100°. 
A  liquid  boils  when  bubbles  of  its  vapor  can  exist  beneath  the  sur- 
face. It  is  evident  that  this  can  occur  only  when  the  pressure  of 
the  vapor  is  at  least  as  great  as  that  upon  the  surface  of  the  liquid. 

The  effect  of  pressure  on  the  boiling-point  of  a  liquid  is  marked. 
For  example,  water  boils  at  60°  when  the  pressure  is  14.9  cm.  of 
mercury  and  at  200°  when  the  pressure  is  1169  cm. 

The  boiling-point  of  a  substance  is  a  characteristic  property 
which  is  often  used  in  identifying  it.  Wood  alcohol  boils  at  66° 
and  grain  alcohol  at  78°;  it  is,  thus,  possible  to  distinguish  these 
substances  from  each  other  by  determining  their  boiling-points. 

179.  The  kinetic  theory  of  gases  helps  materially  in  picturing 
what  happens  when  a  liquid  evaporates  and  boils.     The  molecules 
in  the  liquid  are  in  constant  motion  and  some  of  those  at  the  sur- 
face have  enough  energy  due  to  this  motion  to  overcome  the 
attraction  of  the  molecules  near  them  and  leave  the  liquid.     They 
pass  into  the  space  above,  forming  the  vapor  over  the  liquid,  and 
diffuse  into  the  air.     Some  of  the  molecules  in  the  vapor,  as  the 
result  of  their  constant  motion,  strike  the  liquid,  are  attracted 
by  the  molecules  there,  and  become  again  a  part  of  it.     There  is 
thus  a  constant  interchange,  molecules  passing  from  the  liquid  to 
the  vapor  and  then  back  again.     If  the  vessel  containing  the 
liquid  is  open  to  the  air,  the  molecules  diffuse  into  it  and  are  carried 
away  by  air  currents;   the  liquid  continues  to  lose  molecules  and 
finally  passes  completely  into  the  gaseous  condition  as  the  result 
of  this  process,  which  is  called  evaporation.     If  the  vessel  contain- 
ing the  liquid  is  closed,  the  molecules  in  the  form  of  vapor  cannot 
escape.     When  the  number  of  molecules  which  pass  from  the 
liquid  to  the  vapor  is  equal  to  the  number  which  pass  from  the 


PROPERTIES  OF  GASES,  LIQUIDS,  AND  SOLIDS          161 

vapor  to  the  liquid,  the  two  forms  are  said  to  be  in  equilibrium. 
It  is  important  to  note  that  the  equilibrium  is  dynamic,  not  static; 
that  is  to  say,  it  is  attained  as  the  result  of  equal  movement  in 
opposite  directions — a  molecule  may  be  in  the  vapor  one  instant 
and  in  the  liquid  another.  A  static  equilibrium  would  result  if 
certain  molecules  passed  to  the  vapor  and  remained  there;  at 
equilibrium  there  would  be  no  interchange  between  the  liquid  and 
the  gas. 

When  the  temperature  of  a  liquid  is  raised,  the  heat  absorbed 
increases  the  energy  of  motion — the  kinetic  energy  of  the  mole- 
cules; as  a  result,  they  reach  the  surface  with  greater  force  and  a 
greater  number  pass  into  the  form  of  vapor.  This  is  in  accord 
with  the  fact  that  the  rate  at  which  evaporation  takes  place 
increases  with  rise  in  temperature.  Since  the  pressure  exerted 
by  a  vapor  is  the  result  of  the  impact  of  the  molecules  on  the  walls 
of  the  vessel  which  contains  it,  increase  in  temperature  should  lead 
to  increase  in  vapor-pressure,  a  fact  which  has  already  been  noted. 

The  pressure  on  a  liquid  exposed  to  air  is  made  up  of  the  pres- 
sure of  its  vapor  and  of  the  air.  As  the  temperature  of  the  liquid 
rises,  the  vapor-pressure  increases  as  we  have  seen;  the  molecules 
which  leave  the  surface  increase  in  numbers  and,  finally,  the  space 
above  the  liquid  contains  only  molecules  of  vapor.  At  the  tem- 
perature at  which  this  occurs  the  pressure  of  the  vapor  equals 
that  of  the  air;  bubbles  of  vapor  can  form  in  the  liquid,  and,  as  a 
result,  the  liquid  boils.  If  more  heat  is  now  applied,  the  liquid 
passes  freely  into  vapor,  and  the  temperature  of  the  liquid  remains 
constant.  The  heat  energy  is  used  in  overcoming  the  attraction 
between  the  molecules  of  the  liquid  and  in  producing  a  gas  against 
the  pressure  of  the  atmosphere. 

If  the  pressure  of  a  gas  on  a  liquid  is  increased,  more  molecules 
press  upon  its  surface.  To  overcome  this  pressure  a  greater  num- 
ber of  molecules  of  vapor  must  be  produced  from  the  liquid.  This 
is  accomplished  by  heating  it  to  a  higher  temperature;  the  boiling- 
point,  thus,  increases  with  increase  in  pressure. 

180.  Heat  of  Vaporization. — When  a  liquid  is  heated  the  tem- 
perature rises  until  the  boiling-point  is  reached;  it  then  stays 
constant  as  the  liquid  passes  into  vapor,  although  heat  is  supplied 
to  it.  The  heat  energy  is  used  in  changing  the  liquid  to  a 
gas,  and  no  longer  shows  its  presence  by  an  increase  in  tempera- 


162 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


ture;  for  this  reason  the  heat  absorbed  is  said  to  be  latent.  The 
number  of  calories  required  to  change  1  gram  of  a  liquid  to  vapor  is 
called  the  heat  of  vaporization;  the  value  varies  with  the  tem- 
perature. The  heat  of  vaporization  for  water  at  its  boiling-point  is 
535.9  calories  and  that  for  alcohol  is  206.4  calories. 

An  important  practical  application  of  the  heat  of  vaporization 
is  found  in  ice-machines.  (Fig.  18.)  Ammonia  gas  is  condensed  by 
pressure  into  a  liquid  that  boils  at  —33°,  and  the  heat  of  vaporiza- 
tion of  which  is  330  calories.  When  the  liquid  is  allowed  to  boil, 


Water 


Condenser ••• ~. 


FIG.  18. 


330  calories  are  absorbed  by  each  gram  of  ammonia  as  it  passes 
into  the  gaseous  form.  This  heat  is  extracted  from  the  material 
which  surrounds  the  vessel  in  which  the  ammonia  boils.  In  re- 
frigerating machines  the  evaporation  of  the  liquid  ammonia  takes 
place  in  a  coil  surrounded  by  a  30  per  cent  solution  of  calcium 
chloride,  which  does  not  freeze  at  the  low  temperature  produced. 
The  cold  solution  of  calcium  chloride  is  circulated  around  tanks 
containing  water,  if  ice  is  to  be  made,  or  through  pipes  in  rooms  to 
be  used  for  cold  storage.  The  ammonia  is  pumped  from  the  coil 
into  pipes  where  it  is  again  liquefied  by  pressure,  and  the  heat 
given  off  is  taken  up  by  cold  water,  which  flows  over  the  pipes. 
It  is  next  returned  in  the  liquid  condition  to  the  coil  and  the  process 


PROPERTIES  OF  GASES,  LIQUIDS,  AND  SOLIDS          163 

repeated.  The  heat  absorbed  when  the  liquid  ammonia  boils 
equals  the  heat  set  free  when  the  gas  is  liquefied  under  pressure. 
The  conversion  of  a  gas  into  a  liquid  is  the  process  opposite  to 
that  of  vaporization;  it  is  called  condensation.  In  this  case  heat  is 
set  free,  the  amount  being  equal  to  that  absorbed  in  vaporization. 
This  fact  can  be  represented  in  the  case  of  water  thus : 

H2O  (liquid)  +  9648  cal.  ?=>  H2O  (vapor) 

This  expression  means  that  18  grams  of  water  as  liquid  absorb  9648 
calories  and  pass  into  18  grams  of  water  as  vapor  (steam),  or,  read 
in  the  opposite  direction,  18  grams  of  steam  condense  to  water  and 
liberate  9648  calories.  The  figures  refer  to  the  change  at  100°. 
A  practical  application  of  the  heat  of  condensation  is  made  in 
heating  houses  with  steam.  In  the  boiler-room  the  energy  pro- 
duced by  the  burning  coal  is  taken  up  as  heat  of  vaporization  by 
the  water,  the  steam  passes  through  pipes  to  the  place  to  be 
heated,  and  there  sets  free  the  heat  through  condensation. 

181.  Distillation. — The  combination  of  vaporization  and  con- 
densation is  much  used  in  purifying  liquids.  The  process,  called 
distillation,  consists  in  converting  the  liquid  to  a  gas  and  then 
condensing  it  to  a  liquid.  An  apparatus  for  effecting  these  changes 
is  illustrated  in  Fig.  14,  p.  97.  In  the  flask  the  liquid  is  heated 
to  boiling,  the  vapor  passes  into  the  condenser,  which  consists  of  a 
tube  surrounded  by  an  outer  jacket  through  which  water  flows. 
The  condensed  liquid  is  collected  in  the  receiver.  Every  pure 
liquid  which  boils  without  decomposition  has  a  definite  boiling- 
point.  Since  the  boiling-points  of  different  substances  are  dif- 
ferent, it  is  possible  to  separate  them  by  distillation.  When  a 
mixture  of  two  liquids  is  heated  the  more  volatile  one — that  is, 
the  one  with  the  lower  boiling-point — vaporizes  more  rapidly 
than  the  other.  If  the  vapor  formed  is  condensed  it  will  be  found 
that  the  part  that  distills  first  will  contain  a  much  higher  per- 
centage of  the  low-boiling  liquid  than  of  the  one  with  the  higher 
boiling-point;  whereas  the  part  which  condenses  last  will  be  richer 
in  the  latter.  By  subjecting  the  various  parts  of  the  distillate — 
the  part  condensed — to  further  distillation  the  two  liquids  can,  in 
most  cases,  be  separated.  It  will  be  recalled  that  this  process  is 
used  in  separating  oxygen  and  nitrogen  from  liquid  air  (22). 


164  INORGANIC  CHEMISTRY  FOR  COLLEGES 

182.  Liquefaction  of  Gases. — We  have  seen  what  happens  when 
a  liquid  is  in  equilibrium  with  its  vapor  at  its  boiling-point;    the 
number  of  molecules  that  pass  from  the  liquid  to  the  vapor  equals 
the  number  which  pass  from  the  vapor  to  the  liquid ;  vaporization 
and  condensation  are  both  taking  place  simultaneously  and  to  the 
same  extent.     The  slightest  change  in  temperature  or  pressure  will 
bring  about  a  marked  effect  on  the  equilibrium.     If  we  attempt  to 
raise  the  temperature  of  the  liquid  by  supplying  heat  to  it  and  keep 
the  pressure  constant,  heat  will  be  absorbed  and  the  evaporation 
will  take  place  more  rapidly  than  condensation;   as  a  result  the 
liquid  will  all  change  to  vapor.     If  we  attempt  to  lower  the  tem- 
perature of  the  liquid  the  evaporation  will  take  place  more  slowly 
than  condensation,  and  all  the  vapor  will  finally  be  converted  into 
liquid. 

Change  in  pressure  also  brings  about  evaporation  if  the  tem- 
perature is  kept  constant.  If  the  pressure  is  reduced  complete 
vaporization  takes  place;  if  it  is  increased  the  vapor  changes  to  a 
liquid.  The  changes  of  conditions  at  the  boiling-point  which 
determine  whether  vaporization  or  liquefaction  take  place  are 
exceedingly  small — so  small  that  the  boiling-point  and  liquefaction- 
point  are  practically  the  same.  With  these  facts  in  mind  the  prin- 
ciples underlying  the  liquefaction  of  gases  by  pressure  can  be 
understood.  Take  the  case  of  chlorine.  Under  a  pressure  of  1 
atmosphere  it  boils  at  —33°.  If  we  confine  some  of  the  gas  in  a 
cylinder  provided  with  a  movable  piston  and  by  means  of  the  latter 
compress  it,  a  point  will  be  finally  reached  when  the  gas  turns  to  a 
liquid.  As  the  pressure  increases  the  boiling-point  (liquefaction- 
point)  rises  from  —33°;  when  this  reaches  the  temperature  of 
the  gas  in  the  apparatus,  any  increased  pressure  causes  the  gas  to 
liquefy  in  the  way  explained  above.  If  the  temperature  is  20° 
liquefaction  takes  place  at  6.62  atmospheres.  In  Faraday's 
original  experiment  on  liquefying  chlorine  (124),  the  pressure 
was  obtained  by  heating  the  solid  hydrate  in  a  closed  vessel;  as 
the  compound  decomposed,  more  and  more  of  the  gas  came  off 
and  the  pressure  increased. 

183.  Critical  Temperature. — We  have  seen  that  the  boiling- 
point  of  a  liquid  depends  upon  the  pressure  exerted  upon  it.     A 
study  of  the  effect  of  high  pressures  on  the  boiling-points  of 
liquids   has   led   to   the   discovery   of   a   very   important   fact, 


PROPERTIES  OF  GASES,  LIQUIDS,  AND  SOLIDS          165 

which  will  be  clear  from  the  following  considerations.  The 
boiling-point  of  water  rises  with  increased  pressure  up  to  195 
atmospheres,  where  it  boils  at  358°.  When  pressures  greater 
than  195  atmospheres  are  applied,  water  changes  into  a  gas 
at  358°,  however  great  the  pressure  may  be.  Above  this  tem- 
perature water  can  exist  only  as  a  gas.  The  boiling-points  of  all 
liquids  change  with  pressure,  and  for  each  there  is  a  temperature 
above  which  they  can  exist  as  a  gas  only;  this  temperature  is  a 
characteristic  property  of  liquids  and  is  called  the  critical  tempera- 
ture. The  pressure  which  a  liquid  exerts  at  its  critical  temperature 
is  its  critical  pressure.  These  constants  for  water  are,  as  has  been 
indicated  above,  195  atmospheres  and  358°.  The  critical  tem- 
peratures for  a  few  substances  are  as  follows:  Sulphur  dioxide  156°, 
carbon  dioxide  31.4°,  chlorine  146°,  oxygen  —118°,  hydrogen 
—234°.  It  will  be  seen  that  the  first  three  gases  can  exist  as  liquids 
at  ordinary  temperatures,  but  that  oxygen  remains  as  a  gas,  what- 
ever the  pressure  upon  it,  at  all  temperatures  above  —118°.  In 
his  experiments  Faraday  found  that  certain  gases  resisted  his 
efforts  to  liquefy  them;  he  called  these  permanent  gases.  It  was 
only  when  it  was  discovered  that  for  each  gas  there  was  a  tem- 
perature above  which  it  could  not  exist  in  the  liquid  condition — 
the  critical  temperature — that  Faraday's  unsuccessful  experiments 
with  the  so-called  permanent  gases  could  be  explained.  In  all 
cases  he  had  not  cooled  them  to  their  critical  temperatures,  and  the 
high  pressures  used  did  not  accordingly  lead  to  liquefaction. 

184.  Liquefaction  of  Gases  Possessing  Low  Critical  Tempera- 
tures.— In  order  to  liquefy  a  gas  the  critical  temperature  of  which  is 
very  low,  such  as  oxygen,  use  is  made  of  another  important  principle. 
It  is  necessary  to  do  work  upon  a  gas  to  compress  it;  a  part  of  the 
energy  used  is  transformed  into  heat.  This  fact  is  familiar  to  one 
who  has  used  a  pump  for  inflating  bicycle  tires;  it  requires  work 
to  compress  the  air  and  the  pump  gets  warm.  A  large  part  of  the 
heat  produced  results  from  the  compression  of  the  gas  rather 
than  from  friction  of  the  moving  parts  of  the  pump.  If  the  heat 
generated  on  compressing  a  gas  escapes  and  the  latter  is  then 
allowed  to  expand  to  its  original  volume  and  pressure,  it  will  now 
contain  less  heat  energy  and  its  temperature  will,  as  a  result, 
be  lower  than  it  was  at  first.  This  phenomenon  is  observed  when 
the  compressed  air  in  an  automobile  tire  is  allowed  to  escape  through 


166  INORGANIC  CHEMISTRY  FOR  COLLEGES 

the  valve  of  the  tire.  A  gas,  accordingly,  increases  in  temperature 
on  compression  and  decreases  in  temperature  on  expansion. 
Through  the  application  of  this  principle  it  is  possible  to  lower  the 
temperature  of  ah-  to  a  point  where  its  constituents  liquefy. 

A  number  of  forms  of  apparatus  have  been  devised  for  this  pur- 
pose. In  one  of  these,  air,  after  being  freed  from  moisture  and 
carbon  dioxide,  is  compressed  to  200  atmospheres  by  a  pump  and 
cooled;  it  is  then  sent  through  a  coil  and  allowed  to  expand  through 
a  small  opening  into  a  second  coil  surrounding  the  first.  From  the 
second  coil,  in  which  the  pressure  is  kept  at  about  15  atmospheres, 
the  air  passes  through  a  second  opening  out  of  the  machine.  When 
the  air  expands  from  200  to  15  atmospheres  its  temperature  is 
lowered.  The  gas  cooled  in  this  way  in  passing  through  the  second 
coil,  circulates  around  that  in  the  first,  and  takes  up  a  part  of  its 
heat.  As  a  result,  in  a  short  time  the  incoming  gas  is  colder  than  at 
first,  and  when  it  expands  its  temperature  drops  lower  than  before. 
In  this  way  the  temperature  of  the  gas  at  200  atmospheres  is  con- 
tinually lowered  and,  finally,  when  it  expands  the  temperature 
drops  to  that  at  which  the  gas  is  a  liquid.  Liquid  air  obtained 
in  this  way  boils  at  about  —190°. 

In  liquefying  gases  the  boiling-points  of  which  are  very  low, 
another  principle  is  used.  The  boiling-point  of  a  substance  is 
lowered  by  decreasing  the  pressure  upon  it.  Hydrogen  boils  at 
—252.5°  at  1  atmosphere;  if  it  is  boiled  under  diminished  pressure, 
which  can  be  obtained  by  removing  the  gas  as  quickly  as  it  is 
formed,  the  boiling-point  is  reduced  to  such  a  temperature  that 
helium,  which  boils  at  —268.5°,  can  be  liquefied. 

185.  Solubility  of  Gases  in  Liquids;  Henry's  Law. — Gases 
differ  markedly  in  their  solubility  in  liquids;  in  general,  of  two 
gases  the  one  that  is  more  readily  liquefied  is  the  more  soluble, 
provided  neither  forms  a  compound  with  the  solvent. 

The  solubility  of  a  gas  in  a  liquid  increases  as  the  pressure  on 
the  gas  increases.  Henry  studied  the  effect  of  pressure  on  the  sol- 
ubility of  gases  in  liquids  and  summarized  the  facts  in  the  form 
of  a  law  which  is  known  by  his  name.  The  law  of  Henry  states 
that  the  solubility  of  a  gas  in  a  liquid  is  proportional  to  the  pressure 
of  the  gas.  This  law  expresses  the  behavior  of  gases  only  when 
they  are  under  a  pressure  which  is  far  removed  from  that  at  which 
they  liquefy — that  is,  when  their  volumes  change  with  pressure 


PROPERTIES  OF  GASES,  LIQUIDS,  AND  SOLIDS          167 

in  accordance  with  Boyle's  law.  The  weight  of  carbon  dioxide 
which  dissolves  in  water  at  15°  and  at  2,  3,  and  4  atmospheres 
pressure  is,  respectively,  very  nearly  2,  3,  and  4  times  the  weight 
which  dissolves  at  1  atmosphere  pressure. 

186.  Sublimation. — Solids,  as  well  as  liquids,  change  directly 
to  a  gas,  and  exert  a  vapor-pressure.     In  most  cases  the  pressure  of 
the  vapor  at  ordinary  temperatures  is  very  small,  but  in  many  it 
is  great  enough  to  be  recognized  by  the  fact  that  the  solid  possesses 
a  characteristic  odor.     For  example,  camphor,  iodine,  and  naph- 
thalene, which  is  used  in  "  moth-balls,"  all  give  off  at  room  tem- 
perature appreciable  quantities  of  vapor.     By  cooling  the  vapor  in 
each  of  these  cases  it  can  be  changed  directly  to  the  solid.     The 
transformation  of  a  solid  to  a  gas  and  back  to  the  original  solid 
without  passing  through  the  liquid  state,  is  called  sublimation.     A 
number  of  solids  do  not  melt  when  they  are  heated,  because  the 
vapor-pressure  of  the  gaseous  form  produced  equals  the  pressure 
of  the  atmosphere  at  a  temperature  below  the  melting-point  of 
the  solid.     Application  is  made  of  this  principle  in  the  purification 
of  iodine,  ammonium  chloride,  naphthalene,  and  other  substances 
of  commercial  importance. 

187.  The  Melting-point. — The  change  of  a  liquid  to  a  gas  and 
that  of  a  solid  to  a  gas  have  been  discussed  above;  we  shall  now 
consider  the  relation  between  the  liquid  and  the  solid  state.     The 
temperature  at  which  a  crystalline  solid  changes  to  a  liquid  is 
called  its  melting-point,  and  that  at  which  a  liquid  changes  to  a 
solid  is  called  its  freezing-point.     For  any  one  substance  these 
points  are  the  same.     If  a  solid  is  heated  it  rises  in  temperature 
until  a  point  is  reached  where  it  begins  to  melt ;  this  temperature  is 
called  its  melting-point.     Any  additional  heat  applied  does  not 
affect  the  temperature,  but  transforms  more  of  the  solid  into  liquid. 
The  number  of  calories  required  to  convert  1  gram  of  a  solid  into  a 
liquid  is  called  its  heat  of  fusion.     The  heat  of  fusion  of  ice  is 
79  calories.     If  a  liquid  is  cooled  its  temperature  falls  to  the  freez- 
ing-point, and  stays  constant  until  all  the  liquid  has  changed  to 
solid.     The  heat  lost  in  solidifying  is  equal  in  amount  to  that 
absorbed  in  liquefaction. 

Pressure  has,  as  we  have  seen,  a  marked  effect  on  the  boiling- 
point  of  a  liquid;  it  has  but  little  effect,  however,  on  the  freezing- 
point.  This  is  due  to  the  fact  that  there  is  a  great  change  in 
volume  when  a  liquid  is  converted  into  a  gas,  and  the  latter  in 


168  INORGANIC  CHEMISTRY  FOR  COLLEGES 

forming  must  overcome  the  pressure  upon  it.  In  the  change  from 
the  solid  to  liquid  state  or  the  reverse,  the  volume  change  is  very 
small.  The  boiling-point  of  water  is  raised  about  20  degrees  by 
increasing  the  pressure  upon  it  from  1  to  2  atmospheres.  The 
same  increase  brings  about  a  lowering  of  only  0.008  degree  in  the 
freezing-point.  The  boiling-point  of  a  liquid  always  rises  with  in- 
creased pressure,  because  the  volume  of  the  gas  is  always  greater 
than  that  of  the  liquid.  The  melting-point  of  a  substance  may  be 
raised  or  lowered  by  increasing  the  pressure;  if  the  volume  of  the 
liquid  is  greater  than  that  of  the  solid,  increase  in  pressure  raises 
the  melting-point;  if  the  volume  of  the  liquid  is  less  than  that  of 
the  solid,  increase  in  pressure  lowers  the  melting-point. 

The  use  of  ice  as  a  refrigerating  agent  is  based  on  the  fact  that 
in  melting  it  absorbs  heat,  and  thus  cools  the  surrounding  air. 
When  liquid  air  was  first  made  it  was  thought  that  it  would  be  a 
valuable  cooling  agent  to  replace  ice  on  account  of  its  very  low 
boiling-point.  A  consideration  of  the  heat  changes  involved  when 
liquid  air  boils  and  ice  melts  will  bring  out  the  fact  that  the  former 
is  not  so  good  an  agent  as  ice.  When  1  gram  of  liquid  air  at 
—  190°  changes  to  a  gas  at  this  temperature,  about  50  calories  are 
absorbed.  To  raise  this  air  to  0°  requires  18  calories.  One  gram 
of  liquid  air  will  thus  absorb  68  calories  in  changing  to  air  at  0°. 
At  this  temperature  1  gram  of  ice  absorbs  79  calories  when  it 
melts;  it  is  evident,  therefore,  that  ice  is  more  efficient  in  absorbing 
heat  than  liquid  air.  Water  is  characterized  by  the  fact  that  it 
requires  such  large  amounts  of  heat  to  change  it  from  one  state  to 
another.  This  fact  makes  ice  an  efficient  refrigerating  agent. 

188.  Transition-points. — When  a  substance  is  at  its  melting- 
point  the  solid  and  liquid  forms  of  it  exist,  side  by  side,  in  equilib- 
rium, if  heat  is  neither  added  nor  taken  away;  the  solid  is  melting 
and  the  liquid  is  freezing,  the  two  processes  taking  place  at  the 
same  rate.  If  heat  is  now  added  or  taken  away,  one  form  changes 
to  the  other  without  a  change  in  temperature.  The  temperature 
at  which  this  occurs  is  called  a  transition-temperature ;  in  the  case 
cited  it  is  a  melting-point.  The  boiling-point  is  also  a  transition- 
point,  for  at  this  temperature  a  liquid  changes  into  a  gas  as  heat 
is  applied  without  a  change  in  temperature  taking  place.  The 
word  phase  is  often  used  in  describing  the  substances  which  undergo 
changes  of  this  kind;  water  can,  for  example,  exist  in  three  forms 


PROPERTIES  OF  GASES,  LIQUIDS,  AND  SOLIDS          169 

or  phases — gaseous,  liquid,  and  solid.  At  the  melting-point  the 
liquid  and  solid  phases  are  in  equilibrium,  and  at  the  boiling-point 
the  liquid  and  gaseous  phases.  A  phase  is  any  homogeneous 
constituent  of  a  system  made  up  of  either  different  physical  states 
of  the  same  substance  or  of  different  substances.  In  sublimation 
we  have  two  phases,  the  gaseous  and  the  solid.  If  we  had  a  solu- 
tion containing  salt  and  some  solid  salt  in  a  bottle  half  filled  with 
the  liquid,  we  would  have  three  phases,  the  gaseous,  the  solution, 
and  the  solid.  Many  substances  can  exist  in  more  than  three 
phases,  for  example,  sulphur  exists  as  a  gas,  in  two  distinct  liquid 
forms,  and  in  two  solid  crystalline  forms.  The  temperature  at 
which  any  one  of  these  forms  or  phases  is  in  equilibrium  with  any 
other  phase  is  a  transition-point. 

189.  Superheating  and  Supercooling. — If  water  is  carefully 
heated  in  a  polished  dish  of  silver  it  is  possible  to  raise  its  tempera- 
ture above  100°  without  boiling  taking  place.  Globules  of  water 
suspended  in  oil  have  been  heated  to  145°  without  passing  into 
steam.  The  boiling-point  of  a  liquid  is  not  the  highest  tempera- 
ture to  which  it  can  be  heated,  but  the  temperature  at  which  the 
vapor  and  the  liquid  are  in  equilibrium.  Water  at  a  higher  tem- 
perature than  this  is  said  to  be  superheated.  If  a  trace  of  the 
gaseous  phase  is  introduced  into  superheated  water,  the  change  to 
steam  takes  place  at  times  with  explosive  violence,  and  the  tem- 
perature drops  to  100°.  It  is  evident,  therefore,  that  water  can- 
not be  superheated  in  the  presence  of  steam;  and  if  it  is  desired 
to  raise  it  above  100°  care  must  be  taken  to  prevent  a  bubble  from 
forming.  This  is  accomplished  by  using  a  polished  vessel  free 
from  inequalities  upon  which  a  bubble  can  form.  In  order  to 
prevent  superheating  and  "  bumping  "  when  boiling  a  liquid  in  a 
glass  flask,  pieces  of  broken  porcelain  or  glass  are  introduced  to 
furnish  sharp  points  upon  which  bubbles  of  steam  may  form. 

Liquids  can  also  be  supercooled — that  is,  kept  in  the  liquid 
condition  below  their  freezing-points.  If  a  bit  of  the  solid  phase 
is  introduced  into  the  liquid  the  mass  quickly  freezes;  the  same 
effect  is  produced  by  agitation  or  by  the  introduction  of  a  rough 
surface  upon  which  a  crystal  of  the  solid  can  form. 


170  INORGANIC  CHEMISTRY  FOR  COLLEGES 


EXERCISES 

1.  Upon  what  physical  properties  of  the  substances  involved  are  based 
the  use  of  the  following:    (a)  mercury  in  thermometers,    (6)  a  double  boiler 
in  cooking,   (c)  sulphur  dioxide  as  a  disinfectant,   (d)  steel  in  bridge  construc- 
tion,   (e)  heating  of  houses  with  hot  water,    (/)  wooden  handles  on  soldering 
irons. 

2.  Name  two  conditions  when  it  would  be  advisable  to  use  a  gas  in  a 
thermometer  instead  of  mercury;  give  a  reason  in  each  case. 

3.  Why  is  it  more  economical  to  run  an  engine  with  steam  at  high  pres- 
sure than  at  low  pressure? 

4.  Name  in  each  case  one  use  of  a  substance  which  is  based  in  part  on  the 
following:   (a)  high  density,   (6)  low  density,   (c)  high  specific  heat,   (d]  high 
latent  heat  of  fusion. 

5.  If  a  piece  of  iron  at  room  temperature  is  put  into  100  c.c.  of  hot  water 
what  would  be  the  effect  on  the  temperature  of  the  water?     Why?     If  the 
experiment  were  repeated  using  a  piece  of  gold  of  the  same  weight  and  tem- 
perature would  the  quantitative  result  be  the  same?     If  not,  why? 

6.  What  difference  would  you  observe  if  you  drank  hot  tea  in  one  case 
from  a  cup  made  of  porcelain  and  in  the  other  from  a  cup  made  of  alum- 
inium?    Why? 

7.  Calculate  the  weight  in  grams  of  cubes  of  the  following  substances, 
each  60  cm.  long  on  the  edge:  copper,  lead,  tin.     The  densities  of  the  metals, 
are,  respectively,  8.95,  11.34,  and  7.3. 

8.  A  piece  of  wood  which  measured  2  X  4  X  12  inches  was  found  to 
weigh  2  Ibs.     Calculate-  the  specific  gravity  of  the  wood. 

9.  A  piece  of  apparatus  constructed  of  steel  was  found  to  weigh  750  Ibs. 
What  would  it  weigh  if  it  were  constructed  of  aluminium  and  its  dimensions 
were  the  same?     The  specific  gravities  of  steel  and  aluminium  are  7.6  and 
2.7,  respectively. 

10.  If  equal  weights  of  alcohol  and  water  are  allowed  to  stand  in  open  ves- 
sels which  liquid  would  disappear  first?     Why? 

11.  Certain  substances  which  are  decomposed  when  an  attempt  is  made 
to  distill  them  in  the  type  of  apparatus  represented  in  Fig.  14,  can  be  dis- 
tilled unchanged  if  the  apparatus  is  freed  from  air  before  the  process  is  carried 
out.     Can  you  explain  this  fact? 

12.  If  a  mixture  of  water  and  aniline,  which  is  a  liquid  that  boils  at  182°, 
is  placed  in  the  distilling  apparatus  represented  in  Fig.  14  and  heated  to 
boiling,  and  the  vapor  is  condensed,  the  product  is  a  mixture  of  the  two  liquids. 
Explain  why  aniline  which  boils  at  182°  can  be  "distilled  with  steam"  in  this 
way. 

13.  Would  there  be  any  difference  in  the  cost  of  distilling  1000  gallons  of 
water  and  the  same  volume  of  alcohol?     If  so,  why? 

14.  How  could  you  prepare  relatively  pure  carbon  dioxide  from  the  mix- 
ture of  gases  obtained  when  carbon  is  burned  in  air,  by  using  a  process  based 
on  Henry's  law? 


CHAPTER  XV 
CARBON  AND  ITS  OXIDES 

190.  Carbon  is  one  of  the  most  important  and  interesting  ele- 
ments, because  its  compounds  are  so  widely  distributed  in  nature 
and  play  such  an  important  part  in  daily  life.  All  living  things 
contain  carbon,  and  life  itself  consists  of  a  series  of  chemical  changes 
in  which  carbon  compounds  take  part.  In  organic  chemistry  the 
compounds  of  this  element  are  studied  in  detail;  over  200,000  of 
these  are  known  and  many  have  been  carefully  investigated.  The 
transformations  which  chemists  have  brought  about  in  carbon 
compounds,  have  led  to  the  preparation  of  thousands  of  sub- 
stances which  have  been  put  to  valuable  uses;  these  include  dyes 
of  many  hues,  medicines,  antiseptics,  photographic  developers, 
soaps,  edible  oils,  inks,  perfumes,  and  many  other  substances 
which  add  to  our  comfort  and  pleasure  in  living.  Through  a  study 
of  carbon  compounds  we  learn  what  takes  place  in  the  digestion 
and  assimilation  of  foods;  we  discover  the  difference  in  value  of 
bread  and  meat  in  supporting  life,  and  how  much  of  the  various 
kinds  of  food  a  man  must  eat  to  remain  strong  and  be  able  to  work. 

The  more  the  chemistry  of  the  compounds  of  carbon  is  studied 
the  nearer  we  get  to  an  understanding  of  life  itself — the  greatest 
of  the  unsolved  problems  of  science.  Before  such  interesting 
subjects  can  be  considered,  however,  it  is  necessary  to  study 
carbon  itself  and  the  simpler  compounds  of  the  element.  Carbon 
compounds  are  widely  distributed  in  the  inorganic  world.  Great 
mountain  ranges  are  made  up  of  calcium  carbonate  and  mag- 
nesium carbonate,  and  carbonates  of  other  metals  are  important 
ores  from  which  these  metals  are  extracted  for  industrial  purposes. 
Natural  gas  and  petroleum  are  made  up  of  compounds  of  carbon, 
and  coal  from  which  we  derive  the  energy  to  do  the  world's  mechan- 
ical work,  is  a  mixture  of  carbon  and  some  of  its  derivatives. 

171 


"  172  INORGANIC  CHEMISTRY  FOR  COLLEGES 

191.  Diamond. — Carbon  occurs  in  the  free  condition  as  dia- 
mond, the  chief  sources  being  South  Africa,  Brazil,  and  India. 
Diamonds  are  found  as  crystals,  but  the  crystalline  form  is  more 
or  less  indistinct,  and  they  generally  resemble  rough  pebbles. 
When  perfectly  pure  they  are  colorless,  although  some  that  are 
slightly  colored  by  impurities,  blue,  green,  or  yellow,  are  used  as 
gems.  Black  diamonds  are  used  for  grinding  purposes.  In  order 
to  bring  out  the  brilliancy  of  the  diamond  it  must  be  cut  and  pol- 
ished. Faces  are  ground  upon  it  at  such  angles  that  the  maximum 
amount  of  the  light  which  penetrates  the  diamond  is  reflected 
from  the  rear  faces  back  through  the  stone.  This  light  in  passing 
into  the  air  is  broken  up  into  the  various  colors  of  the  rainbow. 
A  "  brilliant "  is  a  stone  ground  in  the  shape  of  a  many-sided 
pyramid  around  the  edge  of  the  base  of  which  are  cut  faces; 
the  flat  surface  is  the  face  of  the  diamond.  Since  the  diamond  is 
the  hardest  substance  known  it  must  be  polished  with  diamond 
dust. 

Many  diamonds  have  become  famous  on  account  of  their  size, 
or  the  historical  personages  who  owned  them.  The  largest  dia- 
mond known  is  the  so-called  Cullinan;  it  was  found  in  South 
Africa  and  weighed  3032  carats,  which  is  almost  1^  pounds.  This 
priceless  gem  was  presented  to  King  Edward  VII  of  England.  It 
was  cut  into  a  number  of  jewels,  the  largest  of  which  weighs  516 
carats.  The  famous  Kchinoor,  which  is  one  of  the  crown  jewels 
of  England  and  belonged  to  Queen  Victoria,  weighs  106  carats. 
The  international  carat,  which  weighs  200  milligrams,  has  now 
replaced  the  older  English  carat  of  4  grains  (205  mg.).  The  cost 
per  carat  of  a  diamond  increases  rapidly  with  the  size. 

Diamond  is  insoluble  in  all  liquids.  It  is  stable  at  ordinary 
temperatures,  but  when  heated  out  of  contact  with  the  air,  changes 
to  graphite.  At  high  temperatures  it  burns  in  oxygen,  and  carbon 
dioxide  is  formed.  Diamond  has  the  density  3.5,  and  does  not 
conduct  electricity. 

The  preparation  of  diamonds  artificially  has  always  been 
an  interesting  problem  to  the  chemist.  Moissan,  a  French 
chemist,  succeeded  in  1887  in  making  diamonds,  but  the  process 
did  not  yield  a  product  of  commercial  value.  Moissan's  experi- 
ments were  of  great  scientific  interest,  however,  for  they  dem- 
onstrated how  one  allotropic  form  of  carbon  could  be  changed 


CARBON  AND  ITS  OXIDES  173 

to  another,  and  showed  the  relationship  between  amorphous  char- 
coal and  crystalline  diamond.  Charcoal  dissolves  in  about  20 
times  it  weight  of  melted  iron.  When  the  solution  cools  a  part  of 
the  carbon  separates  out  in  crystals  as  graphite,  which  is  a  third 
form  of  carbon.  The  density  of  graphite  is  2.3,  which  is  consider- 
ably less  than  that  of  the  diamond  (3.5).  It  appeared  possible  to 
have  the  carbon  separate  in  the  denser  form  from  its  solution,  if 
great  pressure  were  exerted  upon  it  during  crystallization.  When 
a  mass  of  molten  iron  cools  it  solidifies  first  on  the  outside,  and  a 
solid  shell  is  formed  which  increases  in  thickness  as  the  solidifica- 
tion takes  place.  Iron  expands  on  passing  from  the  liquid  to  the 
solid  condition,  and,  as  a  consequence,  when  the  liquid  core  of  a 
cooling  mass  of  iron  solidifies,  it  exerts  a  tremendous  pressure  on 
anything  in  it,  since  the  solid  shell  prevents  expansion  outward. 
Moissan  found  that  when  carbon  separated  from  melted  iron 
which  was  suddenly  cooled,  it  appeared,  in  part,  as  diamonds. 
These  were  obtained  by  dissolving  the  iron  in  acid,  and  were  found 
to  be  very  small,  the  largest  having  a  diameter  of  about  one-half  a 
millimeter.  It  is  of  interest  to  note  here  that  meterorites  com- 
posed of  metallic  iron  have  been  found  which  contained  graphite 
and  diamonds,  both  black  and  colorless.  The  way  to  prepare 
diamonds  commercially  will  no  doubt  be  discovered.  Artificial 
rubies,  sapphires,  and  other  gems  have  been  made,  which  equal  in 
beauty  those  found  in  nature. 

Diamonds  which  are  of  no  value  as  jewels  on  account  of  their 
color,  are  used  to  cut  glass,  as  tips  for  drills  to  bore  rocks,  and  as 
a  powder  to  polish  jewels.  The  dark-colored  or  black  form  of  the 
diamond,  which  may  occur  in  pieces  weighing  half  a  pound,  is 
called  carbonado  or  bort. 

192.  Graphite. — Graphite  is  a  more  or  less  pure  form  of  car- 
bon, which  occurs  in  crystalline  masses.  It  is  black  or  steel  gray 
in  color,  and  is  made  up  of  flakes  which  are  soft  and  slippery. 
Its  specific  gravity  varies,  that  of  the  purest  varieties  being  about 
2.3.  Graphite  is  widely  distributed  in  small  quantities;  it  is 
mined  for  commercial  purposes  in  New  York,  Canada,  Siberia, 
and  Ceylon.  The  uses  of  graphite  are  based  upon  its  physical 
properties  and  its  chemical  inertness.  It  changes  only  slowly 
when  heated  in  the  air,  and  for  this  reason  is  used  for  making 
crucibles  and  other  apparatus  in  which  chemical  reactions  are  to 


174  INORGANIC  CHEMISTRY  FOR  COLLEGES 

be  carried  out  at  a  high  temperature.  Graphite  conducts  elec- 
tricity and  is  but  slowly  attacked  by  chlorine;  these  properties 
make  it  available  for  electrodes  in  the  electrolytic  industries. 

The  name  given  to  graphite  was  derived  from  the  Greek  word 
meaning  to  write,  on  account  of  the  fact  that  due  to  its  softness  it 
left  a  black  mark  when  drawn  across  paper.  It  was  also  called 
plumbago  or  black  lead,  because  lead  possesses  this  property. 
In  making  lead  pencils  ground  graphite  is  mixed  with  clay  and 
forced  through  dies;  the  sticks  are  then  heated  to  a  high  tempera- 
ture. The  hardness  of  a  pencil  is  determined  by  the  proportion 
of  graphite  mixed  with  the  inert  material.  Graphite  is  used  in 
making  stove  polish,  because  it  protects  the  iron  of  the  stove  from 
the  action  of  oxygen,  and  even  at  high  temperatures  changes  but 
slowly;  and  owing  to  its  crystalline  structure  it  reflects  light  and 
thus  serves  as  a  polish.  The  use  of  graphite  as  a  lubricant  is  based 
upon  the  fact  that  it  is  made  up  of  slippery  scales.  It  can  be  used 
when  the  temperature  is  high,  when  dust  is  apt  to  collect  if  oil  is 
the  lubricant,  and  for  reducing  the  friction  in  machinery  made  of 
wood;  in  the  last  case  oil  would  be  rapidly  absorbed  by  the  wood, 
whereas  graphite  is  not.  On  account  of  the  fact  that  graphite 
conducts  electricity,  it  is  used  to  cover  molds  of  wax  which  are  to 
be  coated  electrolytically  with  metals  in  making  electrotypes. 

A  large  amount  of  graphite  is  now  manufactured  from  coal  at 
Niagara  Falls  by  a  process  invented  by  Acheson.  Ground  anthra- 
cite coal,  with  which  is  mixed  some  iron  oxide  and  sand,  is  placed 
in  a  furnace  made  of  fire-brick.  By  means  of  electrodes  which 
enter  the  furnace  through  the  walls,  a  powerful  alternating  cur- 
rent of  electricity  is  passed  through  the  mixture.  On  account  of 
the  granular  structure  of  the  material  in  the  furnace  heat  is  gen- 
erated as  the  result  of  the  resistance  offered  to  the  current.  Elec- 
tric furnaces  which  are  heated  in  this  way  are  called  resistance 
furnaces;  they  are  much  used  in  the  industrial  preparation  of 
compounds  requiring  very  high  temperatures  for  their  formation. 
When  the  mixture  in  the  furnace  is  heated  in  this  way  for  twenty- 
four  hours,  it  is  changed  to  graphite.  As  the  heat  in  the  furnace  is 
increased  some  of  the  carbon  first  unites  with  the  oxygen  present 
in  the  iron  oxide  and  sand  (silicon  dioxide)  and  carbides  of  iron 
and  silicon  are  formed;  at  the  high  temperature  finally  attained 
the  iron  and  silicon  are  volatilized  from  these  compounds  and  the 


CARBON  AND  ITS  OXIDES  175 

carbon  is  left  as  graphite.  By  molding  the  materials  used  with 
pitch  into  the  particular  forms  desired,  and  then  baking  them  at  a 
high  temperature,  it  is  possible  to  prepare  electrodes,  which  are 
changed  to  graphite  when  heated  in  the  electric  furnace. 

193.  Wood  Charcoal. — Charcoal  is  a  more  or  less  pure  form 
of  carbon,  which  is  obtained  by  decomposing  at  a  high  temperature 
substances  which  contain  compounds  of  carbon.  The  proper- 
ties of  the  resulting  charcoal  are  determined  by  its  source.  Char- 
coal lacks  crystalline  structure;  it  is  said  to  be  amorphous.  The 
purest  variety  is  obtained  by  heating  sugar  to  a  high  temperature 
away  from  the  air.  Larger  quantities  of  charcoal  are  prepared 
from  wood,  on  account  of  the  cheapness  of  the  source.  Logs  of 
wood  are  piled  closely  together  and  covered  with  turf;  they  are 
then  set  on  fire  and  the  amount  of  air  admitted  is  such  that  the 
wood  smolders.  At  the  end  of  a  number  of  days  the  fire  dies  out 
and  the  wood  which  has  not  burned  is  recovered  as  charcoal.  Very 
valuable  products  are  lost  in  this  way,  and  with  the  increasing  value 
of  wood,  charcoal  is  made  more  and  more  by  heating  wood  in  iron 
retorts  so  arranged  that  the  volatile  products  formed  can  be  con- 
densed and  saved.  Since  no  wood  is  burned  the  yield  of  charcoal 
is  greater  than  when  a  kiln  is  used. 

Carbon  is  a  very  inert  element  at  ordinary  temperatures. 
Charcoal  resists  the  action  of  air  and  moisture,  and  for  this  reason 
piles  and  the  ends  of  fence  posts  which  are  to  be  placed  under- 
ground are  often  charred.  It  is  said  that  parts  of  a  bridge  erected 
by  Caesar  over  the  Thames  River  were  discovered  in  modern 
times,  and  the  wood  was  found  to  be  charred. 

Charcoal  is  used  as  a  fuel  in  the  household;  it  is  of  particular 
value  for  broiling  because  it  burns  without  smoke  and  gives  a  high 
heat.  It  is  also  used  in  the  preparation  of  iron  from  its  ores;  the 
carbon  in  this  case  unites  with  the  oxygen  combined  with  the 
metal.  Charcoal  is  one  of  the  ingredients  of  gun-powder  (620). 

Charcoal  retains  the  shape  of  the  wood  from  which  it  was 
prepared.  Since  large  quantities  of  volatile  material  are  lost  in 
converting  wood  to  charcoal,  the  latter  contains  a  mass  of  inter- 
stices and  channels  from  which  these  materials  have  been  driven 
out.  As  a  result,  the  charcoal  is  porous  and  has  a  very  large  sur- 
face—a fact  which  leads  to  important  properties  of  carbon  in  this 
form.  When  gases  come  into  contact  with  solids,  some  of  the 


176  INORGANIC  CHEMISTRY  FOR  COLLEGES 

former  adhere  more  or  less  strongly  to  the  surface  of  the  solid  in  a 
way  that  is  not  as  yet  understood.  If  air  is  passed  through  a  glass 
tube  containing  hydrogen  we  should  expect  that  all  of  the  latter 
would  be  driven  out.  If  this  is  done  at  room  temperature,  and  the 
gas  tested  from  time  to  time,  it  will  be  found  that  no  hydrogen 
can  be  discovered  in  the  gas.  If  now  the  glass  tube  is  heated  to  a 
higher  temperature  and  the  gases  tested  again,  hydrogen  will  be 
found  to  be  present.  This  process  can  be  repeated  until  very 
high  temperatures  are  reached,  and  each  time  increasing  the  tem- 
peratures causes  more  of  the  gas  to  appear.  The  conclusion  is 
drawn  that  in  some  way  the  molecules  of  a  solid  at  its  surface 
attract  the  molecules  of  a  gas  and  hold  them  by  what  is  called 
adsorption.  The  phenomenon  is  one  that  takes  place  at  the  sur- 
face, and  as  a  consequence  charcoal  has  high  adsorptive  powers. 

The  amount  of  gas  held  in  this  way  is  determined  by  the  kind 
of  wood  from  which  the  charcoal  is  made,  as  this  is  a  factor  in  the 
porosity  of  the  charcoal.  The  extent  to  which  a  gas  is  adsorbed  is 
also  determined  by  the  nature  of  the  gas;  in  general,  the  more 
readily  a  gas  is  condensed  by  pressure,  the  more  it  is  adsorbed. 
Charcoal  made  from  boxwood  or  cocoanut  shells  shows  a  high 
power  of  adsorption.  Boxwood  charcoal  adsorbs  90  times  its 
volume  of  ammonia,  55  volumes  of  hydrogen  sulphide,  and  9 
volumes  of  oxygen.  Adsorbed  gases  are  more  reactive  than  those 
in  the  usual  condition.  If  charcoal  containing  chlorine  is  brought 
into  hydrogen  the  gases  unite  at  ordinary  temperatures.  If 
dogwood  charcoal,  as  soon  as  it  is  made,  is  powdered,  it  takes  fire 
as  the  result  of  the  fact  that  the  heat  generated  in  the  adsorption 
of  the  air  causes  the  oxygen  and  carbon  to  unite;  charcoal  that 
behaves  in  this  way  is  said  to  be  pyrophoric.  Powdered  charcoal 
made  into  pellets  is  used  medicinally  to  relieve  the  pain  caused  by 
the  pressures  of  gases  in  the  stomach  and  intestines  resulting  from 
indigestion. 

Charcoal  is  used  to  produce  a  very  high  vacuum.  When  it  is  in 
a  vessel  from  which  as  much  air  as  possible  has  been  pumped  out 
in  the  usual  way,  it  adsorbs  a  large  part  of  what  is  left.  Adsorbed 
gases  can  be  removed  unchanged  from  charcoal  by  heating  it  in  a 
vacuum  to  a  high  temperature.  The  high  adsorbing  quality  of 
charcoal  for  gases  was  utilized  in  the  construction  of  gas  masks 
used  in  warfare.  As  the  result  of  extended  experimentation  char- 


CARBON  AND  ITS  OXIDES  177 

coal  was  prepared  in  such  a  form  that  it  possessed  markedly 
increased  adsorbing  power  over  that  which  had  been  previously 
prepared.  This  was  accomplished  by  heating  the  charcoal  to  a 
high  temperature  in  steam,  and  by  adding  to  it  small  quantities  of 
other  substances. 

Charcoal  is  used  as  a  catalyzer  to  facilitate  the  union  of  gases; 
the  explanation  of  its  behavior  here  is  traceable  to  its  power  to 
adsorb  one  or  both  of  the  gases.  In  the  highly  condensed  condi- 
tion in  which  they  exist  when  adsorbed,  they  react  at  ordinary 
temperatures.  Such  facts  as  these  add  great  interest  to  the  study 
of  the  causes  underlying  adsorption,  and  attention  is  now  being 
devoted  to  this  highly  interesting  and  important  phenomenon. 
When  it  is  more  fully  understood  we  shall  be  nearer  an  explana- 
tion of  certain  kinds  of  catalytic  action. 

Charcoal  was  used  as  a  catalyzer  in  preparing  phosgene,  an 
important  war-gas,  which  has  the  formula  COCk.  When  chlorine 
and  carbon  monoxide,  CO,  are  mixed  and  exposed  to  the  sunlight, 
the  gases  combine  slowly.  This  method  was  inapplicable  when 
large  quantities  of  phosgene  had  to  be  prepared.  It  was  found 
that  the  passage  of  the  gases  over  charcoal  brought  about  their 
union  readily,  and  this  method  was  used. 

194.  Animal  Charcoal. — Charcoal  also  adsorbs  solids  and 
liquids.  Coloring  matters  can  often  be  removed  from  solutions  by 
passing  the  latter  over  charcoal.  For  this  purpose  charcoal  made 
by  heating  bones  to  a  high  temperature  is  very  efficient.  The 
charcoal  in  this  case  is  spread  over  the  mineral  matter  of  which 
bones  are  chiefly  made  up — calcium  phosphate.  Charcoal  made 
in  this  way  is  called  bone-black  or  animal  charcoal.  Crude  sugar 
contains  a  brown  substance;  when  it  is  dissolved  in  water  and  the 
solution  is  allowed  to  trickle  over  bone-black  the  substances 
which  impart  a  color  to  it  are  completely  adsorbed.  From  the 
colorless  solution  pure  white  sugar,  as  we  know  it,  is  obtained 
by  evaporation  and  crystallization.  Coloring  matters  like  litmus, 
indigo,  and  those  in  vinegar,  tea,  etc.,  are  readily  adsorbed  by 
charcoal.  In  general,  substances  whose  molecules  are  large  are 
adsorbed.  The  organic  matter  in  drinking  water  is  also  adsorbed 
by  charcoal,  a  fact  which  was  formerly  utilized  in  the  household; 
but  as  the  niters  containing  charcoal  soon  become  clogged  and 
ceased  to  act,  the  use  for  this  purpose  has  been  discontinued. 


178  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Charcoal  which  is  used  industrially  for  decolorization  is  periodi- 
cally heated  to  a  high  temperature  to  restore  its  activity. 

195.  Lampblack. — When  many  substances  containing  a  large 
percentage  of  carbon  are  burned  in  a  limited  supply  of  air,  the  other 
elements  present  unite  with  the  oxygen,  and  carbon  is  left  in  the 
state  of  a  very  fine  powder.     Amorphous  carbon  prepared  in  this 
way  is  called  lampblack;  it  is  familiar  as  the  soot  formed  when  a 
lamp  smokes.     Lampblack  has  a  number  of  important  industrial 
uses  and  is  prepared  from  residues  obtained  in  the  purification 
of  oils  from  petroleum  and  coal-tar.     The  products  formed  as  the 
result  of  incomplete  combustion  are  led  through  a  series  of  cham- 
bers in  which  the  soot  deposits.     Natural  gas  is  also  a  source  of 
lampblack.     It  is  prepared  by  allowing  the  burning  gas  to  fall  on  a 
rotating  cold  iron  plate  from  which  the  carbon  is  scraped  off.     This 
so-called  gas-black  is  used  in  the  manufacture  of  some  forms  of 
black  rubber  for  automobile  tires.     Lampblack  is  used  in  making 
paints,  shoe  polish,  printer's  ink,  India  ink,  and  other  substances 
which  require  a  black  pigment. 

Coke  is  an  impure  form  of  carbon  which  is  obtained  by  heating 
coal.  It  is  of  such  importance  that  it  will  be  treated  in  detail  in 
the  next  chapter. 

196.  Relation  between  the  Allotropic  Forms  of  Carbon. — The 
fact  that  equal  weights  of  diamond  and  charcoal  give  different 
amounts  of  heat  when  burned  has  already  been  noted.     One  gram 
of  diamond  produces  when  burned  7805  calories,  1  gram  of  graphite 
7850,  and  1  gram  of  sugar  charcoal  8040  calories.     Diamond  and 
graphite  exist  as  crystals  and  have  a  smaller  heat  of  combustion 
than  charcoal;   it  is  evident  that  the  orderly  arrangement  of  the 
atoms  takes  place  with  the  evolution  of  heat.     This  arrangement 
leads  to  a  closer  packing  together  of  the  atoms,  for  the  density  of 
the  diamond  is  3.5,  that  of  graphite  about  2.3,  and  that  of  charcoal 
about  1.5.    One  gram  of  carbon  as  charcoal  occupies  over  twice  the 
space  occupied  by  1  gram  as  diamond;  this  does  not  refer  to  the 
volume  of  the  charcoal,  which  includes  air  spaces  in  its  pores,  but 
only  to  that  occupied  by  the  carbon  itself. 

The  difference  in  the  physical  condition  and  energy  content 
of  the  allotropic  forms  of  carbon  leads  to  a  difference  in  chemical 
activity.  Diamond  is  scarcely  attacked  by  the  most  vigorous 
chemical  reagents,  such  as  powerful  oxidizing  agents,  whereas 


CARBON  AND  ITS  OXIDES  179 

graphite  is  slowly,  and  charcoal  more  rapidly  converted  into  sol- 
uble substances  by  a  mixture  of  nitric  acid  and  potassium  chlorate. 

197.  Properties  of  Carbon. — Many  of  the  properties  of  carbon 
have  been  noted  above.     Carbon  does  not  melt;  at  the  tempera- 
ture of  the  electric  arc  it  volatilizes,  and  even  at  lower  tempera- 
tures   sublimes    slowly.      The    black    coating    that   appears    in 
incandescent  electric  bulbs  in  which  a  carbon  filament  is  used,  is  a 
deposit  of  the  element. 

Under  ordinary  conditions  carbon  is  one  of  the  most  inert  of 
elements,  but  as  the  temperature  is  raised  its  activity  increases 
and,  finally,  at  the  temperature  of  the  electric  arc  it  is  one  of  the 
most  active  of  elements.  When  heated  sufficiently  it  unites  with 
sulphur,  silicon,  iron,  aluminium,  calcium,  and  other  elements; 
the  resulting  compounds  are  called  carbides.  Many  of  these 
which  are  now  of  great  industrial  importance  were  first  made  by 
Moissan,  who  applied  electricity  to  the  production  of  high  tem- 
peratures, and  invented  methods  of  studying  chemical  change  at 
3000°  and  over.  The  more  important  carbides  will  be  described 
later. 

CARBON  DIOXIDE 

198.  On  account  of  the  fact  that  carbon  dioxide  is  present  in 
the  air  and  plays  such  an  important  part  in  animal  and  vegetable 
life,  it  is  a  compound  of  the  first  importance.     It  was  one  of  the 
first  gases  to  be  recognized  as  being  distinct  from  air.     Van  Hel- 
mont  (1577-1644)  studied  the  gas  produced  during  alcoholic  fer- 
mentation and  showed  that  it  was  also  produced  by  the  action  of 
acids  on  chalk  and  when  carbon  burned.     He  called  it   "  gas 
sylvestre."     In  1757,  Joseph  Black,  a  Scotchman,  showed  that 
carbon  dioxide  was  absorbed  by  alkalies  like  sodium  hydroxide, 
and  as  the  gas  disappeared  he  called  it  "  fixed  air."     He  demon- 
strated that  it  was  obtained  when  limestone  was  heated  to  make 
lime,  and  that  an  insoluble  compound  was  formed  when  it  was 
passed  into  limewater,  a  reaction  which  is  used  to-day  as  a  test 
for  the  gas.     Priestley  discovered  the  presence  of  carbon  dioxide 
in  the  air,  and  Lavoisier  showed  that  it  was  formed  in  breathing. 

199.  Occurrence  of  Carbon  Dioxide.— The    gas    is  a  normal 
constituent  of  the  atmosphere  being  present  to  the  extent  of  3 
parts  by  volume  in  10,000.     Its  presence  in  the  air  is  largely  due 


180  INORGANIC  CHEMISTRY  FOR  COLLEGES 

to  the  fact  that  it  is  produced  as  the  result  of  the  decomposition  of 
organic  material,  such  as  decaying  leaves,  trees,  and  other  vege- 
table material;  it  is  also  produced  in  respiration  and  when  carbon 
compounds  burn.  Carbon  dioxide  is  present  in  large  quantities 
in  the  gases  which  issue  from  volcanoes;  it  is  given  off  from  fis- 
sures in  the  earth's  surface,  and  on  account  of  the  fact  that  it  is 
heavier  than  air  it  often  collects  in  caves  and  low-lying  confined 
valleys.  In  the  Grotta  del  Cane  near  Naples  the  bottom  of  the 
cave  is  said  to  be  covered  with  a  layer  of  air  about  18  inches 
deep  containing  a  large  proportion  of  carbon  dioxide.  If  a  dog 
enters  the  cavern  it  is  soon  suffocated,  but  a  man  can  explore 
it  in  safety.  Death  Valley  in  the  Yellowstone  Park  is  so  called 
because  animals  have  died  there  as  a  result  of  the  fact  that  the 
air  was  rendered  unfit  for  breathing  by  carbon  dioxide. 

200.  Preparation  of  Carbon  Dioxide. — Carbon  dioxide  is 
formed  when  carbon  and  any  of  its  compounds  are  burned  in  an 
ample  supply  of  air: 

C  +  O2  =  C02 

If  the  compound  contains  hydrogen,  water  is  also  formed,  for 
example,  the  reaction  which  takes  place  when  methane  burns  is 
as  follows: 

CH4  +  2O2  =  C02  +  2H2O 

Carbon  dioxide  is  used  to  prepare  other  compounds  of  commercial 
importance.  For  this  purpose  the  gas  is  prepared  by  heating 
limestone,  which  is  calcium  carbonate: 

CaCO3  =  CO2  +  CaO 

Calcium  oxide,  which  is  obtained  along  with  carbon  dioxide,  is 
quicklime;  it  was  formerly  called  caustic  lime,  the  word  caustic 
being  derived  from  the  Greek  and  signifying  burnt;  lime  was  pre- 
pared by  "  burning "  limestone.  Other  carbonates  decompose 
into  carbon  dioxide  and  the  oxide  of  the  metal,  or  the  metal  itself, 
when  heated.  The  temperature  at  which  decomposition  takes 
place  freely  varies  with  the  different  carbonates.  It  will  be 
recalled  that  the  less  active  a  metal,  the  more  readily  its  oxide  is 
decomposed  by  heat;  the  same  is  true  of  the  carbonates  although 
these  compounds  are  broken  down  more  readily  than  the  oxides. 
Calcium  oxide  has  not  been  decomposed  by  heat,  but  its  carbonate 


CARBON  AND  ITS  OXIDES  181 

can  be  converted  into  the  oxide  of  the  metal  and  carbon  dioxide; 
at  900°  the  pressure  of  the  gas  given  off  is  equal  to  that  of  the 
atmosphere.  The  carbonates  of  the  active  metals,  sodium  and 
potassium,  decompose  only  very  slowly  at  the  temperature  of  the 
electric  furnace.  When  the  carbonate  of  a  very  inactive  metal 
like  silver  is  heated,  the  temperature  of  decomposition  is  such 
that  the  oxide  cannot  exist,  and,  as  a  consequence,  oxygen  is  also 
lost  and  the  metal  is  left  in  the  free  condition. 

Carbon  dioxide  is  also  prepared  by  the  action  of  acids  on  cal- 
cium carbonate;  this  is  the  method  usually  employed  in  the  labora- 
tory: 

CaC03  +  2HC1  =  CaCl2  +  H2C03 

H2CO3  =  H2O  +  C02 

A  double  decomposition  first  takes  place  as  the  result  of  which 
carbonic  acid,  H^COs,  and  a  salt  of  calcium  are  formed.  The 
former  is  unstable  and  breaks  down  into  carbon  dioxide  and  water. 
The  reaction  takes  place  as  the  result  of  the  fact  that  one  of  the 
products  formed  escapes.  All  carbonates  are  decomposed  by 
acids,  and  the  formation  of  carbon  dioxide  when  they  are  treated 
in  this  way  is  used  as  a  test  for  carbonates. 

Carbon  dioxide  is  formed  when  certain  organic  substances,  like 
the  sugar  present  in  fruits,  undergo  fermentation.  The  gas  formed 
in  this  way  is  collected  and,  stored  under  pressure  in  cylinders,  is 
an  article  of  trade. 

Carbon  dioxide  is  formed  in  the  putrefaction  and  decay  of  all 
organic  substances,  and  most  of  that  which  gets  into  the  air 
comes  from  this  source.  It  is  also  formed  in  respiration. 

201.  Physical  Properties  of  Carbon  Dioxide. — Carbon  dioxide 
is  a  colorless  gas,  which  is  odorless  when  diluted  with  other  gases; 
the  pure  substance,  when  breathed,  produces  a  slight  tingling 
sensation.  It  is  about  1.5  times  as  heavy  as  air;  1  liter  at  760 
mm.  and  0°  weighs  1.976  grams.  On  account  of  this  fact,  the  gas 
can  be  collected  by  upward  displacement  of  air,  and  can  be  poured 
from  one  vessel  to  another  like  water.  Carbon  dioxide  dissolves 
in  water;  the  amount  which  passes  into  solution  is  determined  by 
the  temperature  of  the  water  and  the  pressure  exerted  on  the  gas. 
At  15°  and  760  mm.  pressure  1  liter  of  water  dissolves  1  liter  of  the 
gas;  at  2  atmospheres,  twice  the  weight,  at  3,  three  times;  at  mod- 


182  INORGANIC  CHEMISTRY  FOR  COLLEGES 

erate  pressures  the  solubility  is  proportional  to  the  pressure  of  the 
gas.  A  solution  of  carbon  dioxide  in  water  under  a  pressure  of  3 
to  4  atmospheres  is  sold  as  soda  water.  The  name  is  derived 
from  the  fact  that  baking  soda,  which  is  a  carbonate,  was  formerly 
the  source  of  the  carbon  dioxide  used  in  making  the  water.  When 
a  bottle  containing  soda  water  is  opened  the  pressure  is  reduced  to 
that  of  the  atmosphere,  and  the  carbon  dioxide  comes  off  as  bub- 
bles, until  the  amount  in  solution  is  reduced  to  that  which  water 
can  dissolve  at  the  temperature  of  the  water  and  the  pressure  of 
the  atmosphere. 

202.  As  the  critical  temperature  of  carbon  dioxide  is  31.4°,  it 
can  be  liquefied  by  pressure  alone  at  ordinary  temperatures. 
Liquid  carbon  dioxide  has  the  specific  gravity  0.95  at  0°;  it  has  a 
vapor-pressure  of  59  atmospheres  at  20°  and  must  be  kept  in  cyl- 
inders of  steel  that  can  resist  this  pressure.     When  liquid  carbon 
dioxide  comes  into  the  air  it  evaporates  very  rapidly  and  much 
heat  is  absorbed;    this  heat  is  taken  up,  in  part,  from  the  sub- 
stance itself,  the  temperature  of  which  accordingly  falls,  and  when 
—79°  is  reached  it  solidifies  to  a  white,  snow-like  solid.     It  first 
becomes  a  solid  rather  than  a  liquid  because  at  the  pressure  of 
1  atmosphere  there  is  no  temperature  at  which  gaseous  and  liquid 
carbon  dioxide  are  in  equilibrium;  it  has  no  boiling-point.     If  solid 
carbon  dioxide  at  a  low  temperature  is  heated,  its  vapor-pressure 
increases  until  at  —79°  it  equals  the  pressure  of  the  air,  and,  at 
this  point,  sublimes  without  first  changing  to  the  liquid  state. 
Under  a  pressure  of  3.5  atmospheres  liquid  carbon  dioxide  boils 
at   —56°.      Solid  carbon  dioxide  is  used  as  a  means  of  securing 
low  temperatures  for  experimental  work.     It  is  mixed  with  ether 
or  acetone  in  order  to  bring  the  cooling  agent  in  close  contact  with 
the  vessels  used  to  contain  the  materials  being  studied.     At  —79° 
mercury  is  a  solid,  and  a  bit  of  rubber  tubing  becomes  so  brittle 
it  can  be  broken. 

203.  Chemical  Properties  of  Carbon  Dioxide. — Carbon  dioxide 
is  a  very  stable  compound.     When  heated  to  2000°  it  dissociates  to 
the  extent  of  about  1.8  per  cent  into  carbon  monoxide  and  oxygen: 

2C02  <^  2CO  +  O2 

If  the  temperature  is  lowered  the  products  of  the  dissociation  unite 
to  form  carbon  dioxide.  When  carbon  dioxide  is  led  over  the 


CARBON  AND  ITS  OXIDES  183 

more  active  metals  heated  to  a  high  temperature,  a  decomposition 
like  that  produced  by  heat  takes  place,  but  the  oxygen  unites  with 
the  metal: 

Zn  +  C02  =  ZnO  +  CO 

Since  carbon  dioxide  is  the  product  formed  when  carbon  burns,  the 
gas  does  not  support  the  combustion  of  carbon  or  its  compounds. 
If  a  piece  of  burning  wood  is  put  into  carbon  dioxide  it  is  imme- 
diately extinguished.  Wood  will  not  burn  in  air  which  contains  as 
little  as  15  per  cent  by  volume  of  carbon  dioxide. 

A  solution  of  carbon  dioxide  in  water  shows  the  properties 
which  are  characteristic  of  acids;  it  turns  litmus  from  blue  to  red, 
and  forms  salts  with  bases.  A  reaction  takes  place  between  carbon 
dioxide  and  water  as  the  result  of  which  carbonic  acid  is  formed: 

CO2  +  H2O  <=±  H2C03 

The  reaction  is  a  reversible  one;  as  the  temperature  is  raised  the 
carbonic  acid  dissociates  into  water  and  carbon  dioxide,  which 
escapes  as  a  gas  from  the  solution.  Carbon  dioxide  is  classed  as  an 
acid  anhydride,  because  it  is  formed  as  the  result  of  the  elimina- 
tion of  water  from  an  acid.  Carbonic  acid  reacts  with  sodium 
hydroxide  to  form  sodium  carbonate : 

2NaOH  +  H2C03  =  Na2C03  +  2H2O 

A  different  salt  can  be  formed  by  using  the  amounts  of  the  two 
substances  represented  by  the  following  equation: 

NaOH  +  H2C03  =  NaHC03  +  H2O 

In  this  case  but  one-half  of  the  hydrogen  of  the  acid  is  replaced  by 
sodium;  the  salt  is  called  sodium  bicarbonate.  Acids,  in  general, 
which  contain  2  hydrogen  atoms  can  form  two  kinds  of  salts; 
they  can  react  with  one  molecule  and  with  two  molecules  of  a 
base  like  sodium  hydroxide;  and  for  this  reason  are  called  dibasic 
acids.  Hydrochloric  acid,  HC1,  which  contains  but  1  hydrogen 
atom  is  a  monobasic  acid  and  phosphoric  acid,  H3PO4,  which 
contains  3  atoms  of  hydrogen  that  can  be  replaced  by  metals, 
is  called  a  tribasic  acid. 

The  reaction  which  takes  place  between  carbon  dioxide  and 
sodium  hydroxide  serves  as  a  convenient  means  of  separating  it 


184  INORGANIC  CHEMISTRY  FOR  COLLEGES 

from  other  gases.  If,  for  example,  a  mixture  of  oxygen  and  carbon 
dioxide  is  passed  through  sodium  hydroxide,  the  latter  reacts  and 
sodium  carbonate,  a  solid  soluble  in  the  water,  is  formed;  the 
oxygen  passes  on  unchanged  and  in  this  way  is  freed  from  the  car- 
bon dioxide. 

204.  Carbonates. — The  carbonates  of  many  metals  occur  in 
nature,  some  of  which  are  important  minerals  and  will  be  described 
later.  The  carbonates  belong  to  three  classes — the  normal  car- 
bonates, like  Na2CC>3,  CaCOs,  and  FeCOa,  in  which  no  hydrogen 
is  present,  the  acid  carbonates,  like  NaHCOa,  which  are  formed  as 
the  result  of  the  partial  replacement  of  the  hydrogen  in  carbonic 
acid  by  metallic  atoms,  and  the  basic  carbonates.  To  the  last- 
named  class  belong  the  salts  formed  as  the  result  of  the  partial 
replacement  of  the  hydroxyl  groups  of  bases  by  the  radical  of 
carbonic  acid,  COs;  the  composition  of  basic  salts  is  usually  com- 
plex and  will  be  considered  later.  All  the  normal  carbonates  of 
the  commonly  occurring  elements  are  insoluble  in  water  except 
those  of  sodium,  potassium,  and  ammonium.  It  will  be  well  to 
remember  this  statement,  as  it  summarizes  many  facts  which  will 
be  useful  later.  The  insoluble  carbonates  can,  in  general,  be  made 
by  double  decomposition,  as  illustrated  here  by  the  case  of  calcium 
carbonate : 

CaCl2  +  Na2CO3  =  CaCO3  +  2NaCl 

206.  Test  for  Carbon  Dioxide  and  Carbonates. — When  carbon 
dioxide  comes  in  contact  with  a  solution  of  calcium  hydroxide, 
Ca(OH)2,  calcium  carbonate,  an  insoluble  substance,  is  formed: 

CO2  +  Ca(OH)2  =  CaCO3+  H2O 

This  reaction  serves  as  a  test  for  carbon  dioxide,  because  no  other 
gas  behaves  in  this  way.  The  test  is  made  by  placing  a  glass  rod 
into  a  solution  of  calcium  hydroxide  (lime-water),  and  then  holding 
the  adhering  drop  of  the  solution  in  the  gas.  If  carbon  dioxide  is 
present  the  surface  of  the  drop  will  be  covered  with  a  white  coating. 
If  it  is  desired  to  determine  whether  a  substance  is  a  carbonate,  a 
bit  of  it  is  treated  with  dilute  hydrochloric  acid  in  a  test-tube,  and  a 
rod  moistened  with  lime-water  is  held  in  the  gas  formed.  If  an 
appreciable  quantity  of  the  gas  is  set  free,  it  can  be  poured  after  a 
few  seconds  into  another  tube  about  one-third  filled  with  lime- 


CARBON  AND  ITS  OXIDES 


185 


"Jf»S04 


NaHCOj 

Solution '" 


water,  care  being  taken  to  transfer  the  gas  only.  The  tube  con- 
taining the  lime-water  is  closed  with  the  thumb  and  shaken;  if 
carbon  dioxide  is  present  a  white  cloud  is  produced.  A  solution 
of  barium  hydroxide,  Ba(OH)2,  can  be  used  instead  of  one  of 
calcium  hydroxide;  a  similar  reaction  takes  place  and  barium 
carbonate,  BaCOa,  is  formed. 

206.  Uses  of  Carbon  Dioxide. — The  gas  is  used  in  making 
washing  soda,  baking  soda,  white  lead,  and  other  carbonates,  to 
aerate  water,  and  in  the  preparation  of  sparkling  beverages  like 
ginger  ale,  etc. 

Portable  fire  extinguishers  which  contain  a  solution  of  a  car- 
bonate and  an  acid  are  much  used;  a  common  form  is  sketched  in 
Fig.  19.  The  apparatus  is  nearly  filled 
with  a  solution  of  baking  soda,  and  is 
closed  by  screwing  on  a  cover  to  which  is 
attached  a  frame  containing  a  bottle 
filled  with  concentrated  sulphuric  acid. 
The  neck  of  the  bottle  is  closed  by  a 
loosely  fitting  stopper  of  porcelain  or  glass. 
This  is  provided  to  prevent  the  acid  from 
absorbing  water  from  the  solution;  if  this 
occurs  the  acid  overflows  and  slowly 
reacts  with  the  carbonate.  When  the 
extinguisher  is  to  be  used  it  is  inverted, 
the  stopper  drops  out  of  the  bottle,  and 

the  acid  mixes  with  the  solution  of  the  carbonate.  Carbon  dioxide 
is  formed,  and  rising  through  the  solution,  collects  above  the  liquid. 
The  pressure  produced  forces  the  water  out  of  the  tube  at  the  side. 
A  stream  of  water  containing  sodium  bicarbonate  and  carbon 
dioxide  under  considerable  pressure  is  available,  as  a  result,  for 
extinguishing  fires.  The  fact  that  air  containing  as  little  as  15 
per  cent  of  carbon  dioxide  will  prevent  wood  from  burning  is  one 
of  the  factors  in  making  the  apparatus  very  efficient. 

207.  An  ingenious  method  of  extinguishing  burning  oils  has 
been  recently  developed.     When  water  is  poured  on  a  burning  oil 
the  latter  floats  and  continues  to  burn.     If  water  containing  car- 
bon dioxide  is  used,  the  gas  is  quickly  carried  away  by  the  hot 
products  of  combustion  and  does  not  extinguish  the  fire.     The 
difficulty  was  overcome  by  using  two  solutions  which  were  mixed 


FIG.  19. 


186  INORGANIC  CHEMISTRY  FOR  COLLEGES 

and  allowed  to  flow  over  the  burning  oil.  One  solution  contained 
aluminium  sulphate  and  a  small  amount  of  glue,  and  the  other 
sodium  bicarbonate.  When  the  solutions  are  mixed,  carbon  dioxide 
is  formed,  as  the  aluminium  carbonate  produced  as  the  result  of 
the  double  decomposition  immediately  breaks  down  into  aluminium 
hydroxide  and  carbon  dioxide.  The  glue  in  the  water  serves  to 
retain  the  gas  in  the  form  of  bubbles,  which  persist  for  a  long  time. 
The  foam  produced  in  this  way  covers  the  burning  oil  and  extin- 
guishes it. 

Fire  extinguishers  are  also  provided  which  contain  solid  mate- 
rial to  be  thrown  on  a  flame;  in  these,  a  carbonate  is  used  which 
at  the  temperature  of  the  flame  decomposes  and  liberates  carbon 
dioxide.  The  fact  that  carbon  tetrachloride  is  used  in  one  type 
of  fire  extinguishers  has  already  been  noted;  in  this  case  when  the 
liquid  is  forced  in  a  stream  into  the  fire,  it  changes  to  a  vapor 
which  prevents  combustion. 

208.  Carbon  Dioxide  in  Nature. — The  fact  has  been  mentioned 
that  plants  grow  as  the  result  of  the  transformation  of  water  and 
the  carbon  dioxide  of  the  air  into  the  materials  of  which  they  are 
composed;    and  it  has  also  been  pointed  out  that  when  plants 
decay  they  break  down  into  carbon  dioxide,  which  returns  to  the 
air.     The  changes  which  take  place  are  not  yet  fully  understood 
as  they  occur  step  by  step,  but  the  final  products  of  growth  and 
decay  have  been  carefully  studied.     The  passage  of  carbon  as  car- 
bon dioxide  from  the  air  into  the  materials  of  plant  life,  and 
back  to  the  air  again,  is  sometimes  called  the  cycle  of  carbon. 
Through  this  and  the  cycle  of  other  elements,  such  as  nitrogen, 
the  living  world  constantly  regenerates  itself.     The  plant  is  so 
constituted  that  when  it  dies  the  products  of  its  decomposition 
serve  as  food  for  a  growing  plant.     A  forest  can  regenerate  itself 
forever.     In  the  case  of  animals,  however,  the  waste  materials  or 
the  products  formed  when  death  occurs,  do  not  serve  as  food  for 
other  animals;  they  are  assimilated  by  plants,  however,  and  thus 
return  to  take  part  in  the  cycle  of  carbon.     The  relation  between 
plant  and  animal  life  is  striking;  the  waste  of  the  animal  becomes 
the  food  of  the  plant,  which  in  turn  becomes  the  food  of  the 
animal. 

209.  The  energy  changes  which  take  place  in  the  cycle  of  carbon 
are  of  great  interest.     The  life-processes  of  animals  are  associated 


CARBON  AND  ITS  OXIDES  187 

with  oxidation  in  which  energy  is  set  free  as  heat.  The  foods  we 
eat  pass  into  carbon  dioxide,  which  is  exhaled,  and  heat  is  gene- 
rated in  the  body.  In  the  life  of  an  animal  chemical  energy  is  dis- 
sipated as  heat — the  store  of  available  energy  is  thus  constantly 
decreased.  On  the  other  hand,  when  a  plant  increases  in  weight 
through  growth,  reduction  takes  place  and  energy  is  stored  up. 
When  carbon  dioxide  reacts  with  water  to  form  cellulose,  the  sub- 
stance of  which  the  woody  fiber  of  plants  is  composed,  it  under- 
goes reduction,  and  oxygen  is  set  free.  The  energy  required  is 
taken  up  from  the  sunlight,  for  the  reaction  occurs  only  when  this 
form  of  energy  is  available.  Plants  give  off  carbon  dioxide  at 
night.  Complex  chemical  changes,  which  have  been  studied  care- 
fully but  are  as  yet  not  understood,  occur  near  the  surface  of  the 
leaves  of  plants;  it  is  known,  however,  that  the  green  coloring 
matter  of  the  leaf,  chlorophyll,  plays  an  important  part  in  the 
transformations  as  they  take  place.  If  a  growing  plant  is  placed 
under  water  containing  carbon  dioxide  and  put  in  the  sunlight,  the 
oxygen  generated  can  be  collected  and  tested.  We  see  then,  that 
sunlight  is  the  source  of  the  energy  stored  up  in  plants.  When  we 
climb  a  hill  we  are  using  energy  from  this  source,  when  we  run  an 
engine  by  burning  wood  or  coal,  the  sun  is  doing  the  work.  When 
we  use  a  water-fall  to  generate  electricity  we  are  again  obtaining 
available  energy  which  came  originally  from  the  sun;  for  the  heat 
of  the  sun  evaporated  the  water  from  the  surface  of  the  earth,  and 
the  vapor  rose,  formed  clouds  which  turned  into  rain,  and  finally 
fell  on  an  elevated  part  of  the  earth's  surface,  and  in  seeking  a 
lower  level  the  water  gave  up  a  part  of  its  potential  energy. 

CARBON   MONOXIDE 

210.  When  charcoal  smoulders  a  gas  is  formed,  which  is  char- 
acterized by  the  fact  that  it  is  deadly  poison;  it  is  called  carbon 
monoxide  and  has  the  composition  represented  by  the  formula  CO. 
The  gases  formed  as  the  result  of  firing  explosives  contain  carbon 
monoxide;  as  a  consequence,  the  air  in  the  turrets  of  battleships 
may  contain  the  gas  after  the  guns  have  been  repeatedly  fired. 
Special  precautions  are  taken,  therefore,  to  assure  the  adequate 
ventilation  of  the  turrets.  During  the  recent  war  a  gas  mask 
was  devised  that  furnishes  protection  against  the  gas.  Carbon 


188  INORGANIC  CHEMISTRY  FOR  COLLEGES 

monoxide  was  not  used  as  a  war  gas  on  account  of  its  lightness 
and  its  relatively  low  toxicity. 

Carbon  monoxide  is  an  active  poison.  The  gas  unites  directly 
with  the  red  coloring  matter  of  the  blood,  and  thus  prevents  the 
formation  of  the  compound  of  the  latter  with  oxygen.  It  will  be  re- 
called that  oxygen  is  carried  to  all  parts  of  the  body  by  the  blood ; 
if  this  is  prevented  death  results.  The  compound  formed  when 
carbon  monoxide  gets  into  the  blood  has  a  bright-red  color;  it  is  for 
this  reason  that  the  skin  of  persons  poisoned  by  the  gas  has  such  a 
pink  color  after  death.  Ten  cubic  centimeters  of  carbon  monoxide 
per  kilogram  weight  of  an  animal  will  cause  death;  a  man  of  average 
weight,  will  be  killed  by  inhaling  about  800  c.c.  of  the  gas.  Air 
which  contains  1  part  of  carbon  monoxide  in  800  will  prove  fatal 
in  about  thirty  minutes;  as  little  as  1  part  in  2000  causes  un- 
consciousness and  convulsions,  and,  finally,  death.  The  presence 
of  very  small  amounts  of  carbon  monoxide  in  the  blood  can  be  de- 
tected by  examining  by  means  of  a  spectroscope  (615)  the  light 
which  passes  through  it;  the  blood  absorbs  a  part  of  the  light 
which  passes  through  unchanged  when  pure  blood  is  examined. 
Traces  of  carbon  monoxide  have  been  found  in  tobacco  smoke;  it 
is  claimed  that  the  injurious  effects  of  inhaling  the  gases  produced 
in  smoking  can  be  traced  to  this  cause. 

211.  Preparation  of  Carbon  Monoxide. — When  carbon  dioxide 
is  passed  over  red-hot  zinc  or  iron  it  is  reduced  to  carbon  monoxide 
(203) ;  it  is  also  formed  when  carbon  is  used : 

CO2  +  C  =  2CO 

This  is  a  reaction  of  importance  and  is  used  in  the  preparation  of 
producer  gas,  which  is  described  later  (229).  Carbon  monoxide  is 
formed  along  with  hydrogen  when  steam  is  passed  over  red-hot  coal : 

H2O  +  C  =  CO  +  H2 

This  reaction  is  the  basis  of  the  preparation  of  water-gas,  which  is 
used  for  purposes  of  illumination  and  as  a  source  of  heat  and 
power  (228). 

In  the  laboratory  carbon  monoxide  is  most  conveniently  pre- 
pared by  treating  formic  acid  with  concentrated  sulphuric  acid; 


CARBON  AND  ITS  OXIDES  189 

the  reaction  consists  in  the  removal  of  water  from  formic  acid,  the 
sulphuric  acid  serving  as  the  dehydrating  agent : 

H2CO2  =  H20  +  CO 

The  preparation  is  carried  out  by  allowing  formic  acid  to  fall, 
drop  by  drop,  from  a  funnel  provided  with  a  stop-cock,  into  con- 
centrated sulphuric  acid;  the  latter  is  contained  in  a  flask  con- 
nected with  a  delivery  tube  arranged  to  collect  a  gas  over  water. 
When  oxalic  acid  is  heated  with  sulphuric  acid  both  carbon 
monoxide  and  carbon  dioxide  are  formed: 

H2C2O4  =  H2O  +  CO  +  CO2 

Carbon  monoxide  is  separated  by  passing  the  gas  through  a  sol'i- 
tion  of  sodium  hydroxide;  the  carbon  dioxide  is  converted  into 
sodium  carbonate,  which  remains  in  the  solution,  and  the  carbon 
monoxide  passes  on  unchanged. 

212.  Physical  Properties  of  Carbon  Monoxide. — Carbon  mon- 
oxide is  a  colorless  gas,  with  practically  no  odor  or  taste.     It  is 
slightly  lighter  than  air;  1  liter  at  0°  and  760  mm.  pressure  weighs 
1.250  grams.    Water  dissolves  at  20°  about  2  per  cent  of  its  volume 
of  the  gas.     It  has  been  condensed  to  a  liquid  which  boils  at  — 190° 
and  solidifies  at  -203°. 

213.  Chemical    Properties    of    Carbon    Monoxide. — Carbon 
monoxide  is  not  active  at  ordinary  temperatures.     It  burns  with  a 
blue  flame  when  ignited : 

2CO  +  O2  =  2CO2 

The  blue  flames  observed  over  the  surface  of  coal  that  is  burning 
slowly  under  a  weak  draught  are  produced  as  the  result  of  the 
combustion  of  carbon  monoxide.  Carbon  monoxide  unites  with 
chlorine  in  the  presence  of  sunlight  to  form  carbonyl  chloride  (phos- 
gene) : 

CO  +  C12  =  COC12 

Carbonyl  chloride  boils  at  8°,  and  is  decomposed  by  water  into 
carbon  dioxide  and  hydrochloric  acid;  it  is  used  in  the  preparation 
of  dyes,  and  was  an  important  war-gas. 

At  high  temperatures  carbon  monoxide  is  a  powerful  reducing 
agent;   the  gas  is  formed  in  the  furnaces  used  to  extract  metals 


190  INORGANIC  CHEMISTRY  FOR  COLLEGES 

from  their  ores,  and  is  the  chief  agent  in  effecting  their  reduction. 
The  reactions  with  copper  oxide  and  ferric  oxide  are  represented 
by  the  following  equations : 

CuO  +  CO  =  Cu  +  CO2 
Fe2O3  +  SCO  =  2Fe  +  3CO2 

214.  Test  for  Carbon  Monoxide. — The  fact  that  carbon  monox- 
ide burns  with  a  blue  flame  and  does  not  support  combustion  is 
used  as  a  test  for  the  gas;  this  behavior  is  not  characteristic,  how- 
ever, and  the  conclusion  that  a  gas  is  carbon  monoxide  because  it 
acts  in  this  way  must  be  confirmed  by  other  tests.     Carbon 
monoxide  dissolves  in  a  solution  of  cuprous  chloride,  CuCl,  in 
ammonia.     The  reagent  prepared  in  this  way  is  used  to  separate 
carbon  monoxide  from  other  gases,  and  finds  a  valuable  application 
in  the  analysis  of  flue  gases,  which  contain  carbon  monoxide. 

215.  Carbon  Bisulphide. — When  the  vapor  of  sulphur  is  passed 
over  charcoal  or  coke  at  red  heat,  carbon    disulphide,  CS2,  is 
formed.     The  reaction  is  commonly  carried  out  in  a  furnace  in 
which  an  electric  current  is  the  source  of  heat.     Carbon  disulphide 
is  an  endothermic  compound ;  when  the  liquid  is  formed  the  energy 
change  is  represented  by  the  following  equation: 

C  +  S2  =  CS2  -  19,600  cal. 

For  this  reason  carbon  does  not  burn  in  sulphur  vapor  as  it  does 
in  oxygen.  The  ease  with  which  carbon  disulphide  decomposes, 
and  the  low  temperature  at  which  it  begins  to  burn  (about  200°) 
can  be  traced  to  the  fact  that  the  molecule  contains  more  chemical 
energy  than  the  free  elements  of  which  it  is  composed. 

Carbon  disulphide  is  a  colorless  liquid  which  has  the  specific 
gravity  1.26  at  20°;  when  pure  it  has  an  ethereal  odor,  but  as 
ordinarily  obtained  it  contains  compounds  of  sulphur  which  impart 
to  it  a  very  unpleasant  smell.  Carbon  disulphide  boils  at  46°, 
and  vaporizes  rapidly  at  room  temperature;  the  vapor  is  about  2.7 
times  as  heavy  as  air  and  does  not,  therefore,  diffuse  rapidly. 
Owing  to  this  and  the  fact  that  the  vapor  burns  when  heated 
with  air  to  about  200°,  carbon  disulphide  is  a  dangerous  substance 
to  work  with  in  the  neighborhood  of  flames.  As  in  the  case  of 
other  inflammable  liquids,  a  mixture  of  its  vapor  and  air  explodes 


CARBON  AND  ITS  OXIDES  191 

when  ignited.  When  carbon  disulphide  burns,  carbon  dioxide  and 
sulphur  dioxide,  a  gas  with  a  pungent  disagreeable  odor,  are  formed : 

CS2  +  3O2  =  CO2  +  2SO2 

Carbon  disulphide  is  almost  insoluble  in  water,  but  mixes  in  all 
proportions  with  ether,  benzene,  alcohol,  and  many  oils.  It  is  an 
excellent  solvent;  it  dissolves  phosphorus,  iodine,  sulphur,  wax, 
tars,  resins,  rubber,  oils,  and  fats;  on  this  account  it  is  employed 
as  an  extracting  agent,  but  owing  to  the  danger  connected  with  its 
use  it  is  being  replaced  when  possible  by  carbon  tetrachloride. 
The  vapor  of  carbon  disulphide  is  poisonous,  a  fact  that  leads  to 
its  use  for  exterminating  moles,  rats,  woodchucks,  etc.;  it  is  also 
used  as  a  germicide  and  insecticide. 

216.  Carbon  Tetrachloride. — On    account    of    the    fact    that 
carbon  tetrachloride  is  a  good  solvent  for  oils  and  grease  and  is  non- 
inflammable,  it  has  been  recently  much  used  to  extract  wool,  seeds, 
etc.,  instead  of  gasoline  or  benzine.     A  preparation  to  replace 
gasoline  as  a  cleaning  agent  in  the  household  is  sold  under  the  name 
of  "  carbona  "  which  consists  of  benzine  and  carbon  tetrachloridc 
in  such  proportions  that  the  mixture  or  its  vapor  will  not  burn. 
When  carbon  tetrachloride  is  poured  on  a  fire  the  vapor  produced 
from  it  prevents  combustion.     It  will  be  recalled  that  the  presence 
of  but  15  per  cent  of  carbon  dioxide  in  air  prevents  combustion; 
it  is  probable  that  a  small  proportion  of  the  vapor  o'f  carbon  tetra- 
chloride acts  in  the  same  way.     This  fact  is  utilized  in  the  so-called 
"  pyrene  "  fire  extinguisher,  which  is  of  particular  value  in  extin- 
guishing burning  oils.    If  water  is  applied,  the  oil  floats  on  its  sur- 
face and  continues  to  burn,  but  when  carbon  tetrachloride  is  used 
it  mixes  with  the  oil  and  finally  renders  it  non-inflammable. 

Carbon  tetrachloride  is  made  by  passing  dry  chlorine  into  car- 
bon disulphide  which  contains  a  little  iodine  that  acts  as  a  cata- 
lytic agent: 

CS2  +  3C12  =  CCU  +  S2C12 

The  carbon  tetrachloride  is  distilled  from  the  mixture  and  purified ; 
the  sulphur  chloride  is  recovered  and  used  in  vulcanizing  rubber. 
Carbon  tetrachloride  boils  at  77°  and  has  the  specific  gravity  1.628. 

217.  Calcium  Carbide. — When  lime,  CaO,  is  heated  in  an  elec- 


192  INORGANIC  CHEMISTRY  FOR  COLLEGES 

trie  furnace  with  carbon,  the  oxide  is  reduced  and  the  metal  liber- 
ated unites  with  the  excess  of  carbon  present  to  form  a  carbide : 

CaO  +  3C  =  CaC2  +  CO 

The  preparation  is  carried  out  in  a  furnace  of  the  resistance  type 
(192) ;  it  is  built  of  fire-brick  and  lined  with  carbon,  which  serves 
as  one  electrode.  The  furnace  is  filled  with  coarsely  pulverized 
lime  and  coke,  and  rods  of  carbon  are  introduced  through  the  top  of 
the  furnace  to  serve  as  the  second  electrode.  As  the  calcium  car- 
bide is  formed  it  settles  to  the  bottom  of  the  furnace  as  a  liquid, 
which  is  drawn  off  from  time  to  time. 

Calcium  carbide  as  ordinarily  obtained  is  a  hard,  crystalline 
substance  of  dark  color,  but  when  chemically  pure  is  white.  It 
decomposes  in  the  air  as  the  result  of  the  action  of  the  water- vapor 
present.  With  water  it  decomposes  rapidly  and  calcium  hydroxide 
and  acetylene  are  formed : 

CaC2  +  2H2O  =  C2H2  +  Ca(OH)2 

Commercial  calcium  carbide  is  about  80  per  cent  pure;  it  is  used  to 
make  acetylene  and  in  the  manufacture  of  calcium  cyanamide 
(342). 

218.  Silicon  Carbide. — When  sand,  silicon  dioxide,  is  heated  to 
a  high  temperature  with  carbon,  a  reaction  analogous  to  that 
between  calcium  oxide  and  carbon  takes  place;  in  this  case  silicon 
carbide,  SiC,  and  carbon  monoxide  are  formed: 

SiO2  +  3C  =  SiC  +  2CO 

The  reaction  is  carried  out  in  a  resistance  electric  furnace  like 
that  described  in  connection  with  the  preparation  of  graphite 
(192).  Between  the  ends  of  the  graphite  electrodes  which  enter 
at  the  sides  of  the  furnace,  is  packed  a  core  of  granulated  coke  to 
serve  as  the  conductor  of  the  current.  Around  and  above  this  is 
placed  a  mixture  of  sand,  powdered  coke,  and  salt,  the  latter  being 
added  to  assist  in  binding  the  materials  together.  After  the  reac- 
tion has  taken  place,  the  silicon  carbide  is  found  as  a  layer  of  bril- 
liant black  iridescent  crystals  around  the  central  core  of  carbon, 
which  contains  graphite  and  unchanged  coke.  The  graphite  is 
produced  as  the  result  of  the  dissociation  of  the  silicon  carbide  at 
the  hottest  part  of  the  furnace. 


CARBON  AND  ITS  OXIDES  193 

Silicon  carbide  is  characterized  by  its  hardness,  and  its  use  as  a 
grinding  material  is  based  on  this  fact.  It  has  largely  replaced 
corundum,  or  emery,  for  this  purpose,  and  is  called  in  trade  car- 
borundum. It  is  used  to  make  grindstones,  knife  sharpeners,  etc. 

EXERCISES 

1.  (a)  Name  several  ways  in  which  manganese  dioxide  and  charcoal, 
both  in  the  form  of  a  fine  black  powder,  could  be  distinguished  from  each  other. 
(6)  Which  method  could  be  used  most  conveniently?     (c)  How  could  you 
distinguish  a  mixture  of  equal  weights  of  manganese  dioxide  and  charcoal 
from  pure  charcoal  and  from  pure  manganese  dioxide?     (d)  How  could  you 
separate  charcoal  from  a  mixture  of  it  with  manganese  dioxide? 

2.  (a)  How  could  you  free  air  from  carbon  dioxide?     (6)  How  could  you 
determine  the  percentage  of  carbon  dioxide  in  a  sample  of  air? 

3.  How  could  you  obtain  pure  carbon  dioxide  from  the  mixture  of  nitro- 
gen and  carbon  dioxide  which  results  when  carbon  is  burned  in  air? 

4.  Would  you  expect  calcium  carbonate  or  copper  carbonate  to  be  more 
readily  decomposed  by  heat?     Give  a  reason  for  your  answer. 

5.  Calculate  the  percentage  by  weight  of  carbon  dioxide  in  a  water  solu- 
tion of  the  gas  at  15°  and  760  mm.,  assuming  that  no  carbonic  acid  is  formed. 

6.  What  weight  of  calcium  carbonate  must  be  dissolved  in  an  acid  to 
furnish  10  liters  of  carbon  dioxide  at  0°  and  760  mm.? 

7.  How  much  lime,  CaO,  can  be  prepared  from  1  ton  of  limestone  which 
contains  95  per  cent  CaCO3? 

8.  How  could  you  tell  the  percentage  of  Na2CO3  in  a  mixture  of  Na2CO3 
and  NaCl? 

9.  (a)  What  weight,  approximately,  of  CO2  must  be  mixed  with  the  air 
in  a  room  3X5X6  meters  in  order  to  render  the  air  incapable  of  support- 
ing combustion?     (6)  What  weight  of  washing  soda,  Na2CO3,  10H2O,  would 
have  to  be  treated  with  an  acid  to  produce  this  amount  of  the  gas? 

10.  What  volume  of  carbon  monoxide  and  hydrogen  at  0°  and  760  mm. 
is  produced  when  1  ton  of  coke  containing  90  per  cent  carbon  reacts  with 
steam?     (One  pound  molecular  weight  of  a  gas  occupies  359  cu.  ft.  at  0°  and 
760  mm.) 

11.  What  volume  of  CO  is  obtained  when  50  grams  of  formic  acid  are  de- 
composed by  sulphuric  acid? 

12.  How  could  you  separate  the  CO2  and  CO  formed  by  decomposing 
oxalic  acid  with  sulphuric  acid,  and  obtain  the  two  gases  in  pure  condition? 

13.  Devise  a  method  to  determine  the  amount  of  each  of  the  following 
gases  in  a  mixture :  N2,  O2,  CO,  CO2.     Make  use  of  the  solubilities  of  the  gases 
in  different  reagents. 

14.  What  volume  of  air  contains  just  enough  oxygen  to  react  with  the  gases 
formed  when  12  grams  of  carbon  decompose  water  into  carbon  monoxide 
and  hydrogen? 


CHAPTER  XVI 
COAL,  COKE,  ILLUMINATING  GAS,  FLAMES 

219.  Coal  ranks  first  among  the  natural  resources  which  are 
utilized  by  man  in  producing  the  necessities  and  comforts  of 
modern  civilization.     It  is  the  chief  source  of  heat  and  power,  and 
is  used  in  extracting  metals  from  their  ores.     The  industrial  posi- 
tion of  a  nation  is  determined  largely  by  its  supply  of  coal — a  fact 
that  has  led  to  wars  and  has  influenced  the  political  history  of  the 
world.     Petroleum,  natural  gas,  and  waterfalls  are  also  a  source  of 
power,  but  they  furnish  but  a  very  small  fraction  of  the  energy 
required  to  do  the  world's  work. 

220.  Coal. — Coal  is  the  product  of  the  slow  decomposition  of 
vegetable  material  in  the  absence  of  oxygen.     The  early  stages  of 
this  decomposition  can  be  seen  in  marshes  where  grass,  leaves, 
and  boughs  of  trees  are  undergoing  a  change  which  leads  to  the 
formation  of  peat,  carbon  dioxide,  and  methane,  CEU.     When  a 
stagnant  pool  in  which  these  changes  are  taking  place  is  stirred,  the 
bubbles  of  gas  which  rise  to  the  surface  can  be  lighted  and  burn  as 
the  result  of  the  presence  of  methane,  which  is  also  called  for  this 
reason  marsh  gas.    As  this  change  continues,  more  and  more  of  the 
oxygen  is  removed  from  the  cellulose  of  which  the  vegetable  mate- 
rial is  made  up.     Cellulose  contains  carbon,  hydrogen,  and  oxygen 
in  the  proportions  represented  by  the  formula  CeHioC^,  although 
the  actual  number  of  atoms  of  each  element  present  is  unknown, 
and  for  this  reason  its  formula  is  usually  written  (CeHioOs)*.     The 
removal  of  oxygen  produces  products  which  are  known  successively 
as  peat,  lignite,  brown  coal,  and  bituminous,  semi-bituminous, 
and  anthracite  coal.    The  final  changes  which  lead  to  the  formation 
of  the  various  grades  of  coal  have  been  brought  about,  in  all  prob- 
ability, as  the  result  of  geological  changes  on  the  earth's  surface, 
which  produced  heat  and  pressure  in  the  absence  of  air.     Bitu- 

194 


COAL,  COKE,  ILLUMINATING  GAS,  FLAMES  195 

minous,  or  soft  coal,  contains  carbon,  hydrogen,  and  oxygen 
approximately  in  the  proportion  represented  by  the  formula 
CaeHboC^,  although  it  is  not  a  compound  which  has  this  com- 
position, but  is  a  mixture  of  many  substances.  The  comparison 
of  this  formula  with  that  of  cellulose  will  indicate  that  soft  coal 
contains  a  smaller  proportion  of  hydrogen  and  oxygen  than  the 
substance  from  which  it  was  formed.  Anthracite,  or  hard  coal, 
contains  much  less  oxygen  and  hydrogen  than  soft  coal. 

The  material  from  which  coal  was  formed  contained  in  addi- 
tion to  carbon, hydrogen, and  oxygen,  compounds  in  which  nitrogen, 
sulphur,  iron,  silicon,  and  other  elements  were  present,  and  as  a 
result,  these  elements  are  found  in  coal.  When  coal  is  heated  in 
the  absence  of  air,  more  or  less  volatile  material  is  given  off  which 
contains  ammonia,  NHs,  hydrogen  sulphide,  EkS,  cyanogen, 
C2N2,  and  other  gases,  some  of  which  are  composed  of  carbon  and 
hydrogen,  and  are  called  hydrocarbons.  If  the  residue,  which  is 
called  coke,  is  now  heated  in  the  air  the  carbon  is  burned  and  an 
ash  remains,  which  contains  the  non-volatile,  inorganic  material 
in  the  coal.  The  ash  of  coal  is  a  complex  mixture  which  may  con- 
tain the  silicates  of  calcium,  magnesium,  and  iron,  and  the  oxide 
and  sulphide  of  iron. 

221.  The  Analysis  of  Coal. — The  value  of  a  coal  as  a  source  of 
heat  is  determined  by  its  so-called  calorific  value,  that  is,  the  heat 
produced  when  a  definite  weight  of  it  burns;  and  its  value  for 
making  illuminating  gas  is  determined  by  the  percentage  of  gases 
given  off  when  it  is  heated.  In  analyzing  coal  1  gram  of  the  pow- 
dered air-dried  sample  is  first  heated  at  105°  to  determine  the  free 
water  present.  The  presence  of  water  lessens  the  value  of  the 
coal,  not  only  because  it  will  not  burn,  but  also  because  it  absorbs  a 
large  amount  of  heat  when  it  is  vaporized.  A  crucible  containing 
the  sample  is  covered  and  heated  to  redness  in  the  Bunsen  flame. 
The  loss  in  weight  is  a  measure  of  the  volatile  material  given  off. 
The  coke  that  remains  is  then  heated  in  the  presence  of  air  and 
the  carbon  burned  off,  the  loss  in  weight  being  the  result  of  burning 
the  so-called  fixed  carbon.  The  residue  is  ash.  A  sample  of  the 
coal  is  burned  in  a  bomb-calorimeter  to  determine  its  calorific 
value.  The  heat  produced  is  expressed  for  commercial  purposes  in 
British  thermal  units  (B.t.u.)  per  pound  of  coal.  The  relation 
between  this  unit  and  the  calorie  has  been  explained  in  section  153. 


196 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


Coals  from  different  localities  differ  in  composition.  The 
figures  in  the  following  table  are  the  average  percentages  obtained 
from  analyses  of  a  large  number  of  samples  of  the  fuels  listed. 

ANALYSES  OF  FUELS 


Vola- 

Water. 

tile 

Carbon. 

Ash. 

B.t.u.  per  pound. 

Matter. 

Wood  

20 

49 

30 

1 

7,000  to    9,000 

Peat  

20 

51.6 

25 

3.2 

7,000  to    9,000 

Bituminous  coal  

0.9 

27.4 

64.1 

7.6 

11,000  to  14,000 

Semi-bituminous  coal  .  .  . 

0.5 

16.7 

77.3 

5.5 

13,000  to  15,000 

Anthracite  coal  

2 

4.3 

86.5 

7.2 

12,000  to  14,500 

Coke                         .    . 

2.5 

1.3 

86.3 

12.4 

12,000  to  14,000 

222.  The  Burning  of  Coal. — The  products  formed  when  coal 
burns  are  determined  by  the  proportion  of  air  used ;  if  an  excess  of 
oxygen  is  furnished,  the  carbon  is  converted  into  the  dioxide,  the 
hydrogen  into  water,  the  sulphur  into  sulphur  dioxide,  and  the 
nitrogen  escapes  in  the  elementary  condition.  If  there  is  not 
enough  oxygen  to  form  these  compounds  the  carbon  is  converted 
in  part  into  carbon  monoxide,  and  some  of  it  may  escape  as  smoke; 
and  the  sulphur  may  pass  off  with  the  products  of  combustion  as 
hydrogen  sulphide.  In  burning  coal  to  produce  heat  for  indus- 
trial purposes  attention  is  paid  to  the  manner  in  which  the  coal  is 
added  to  the  furnace — the  stoking — and  to  the  amount  of  air 
admitted  to  the  fire.  If  coal  is  thrown  on  the  fire  so  that  it  forms  a 
thick  layer,  a  part  of  it  is  decomposed  by  the  heat,  and  the  volatile 
matter  produced  is  driven  up  the  chimney  along  with  small  par- 
ticles of  solid,  by  the  hot  products  of  combustion.  There  is,  as  a 
result,  a  loss  of  combustible  matter  and  smoke  is  produced. 

If  not  enough  air  is  admitted  to  the  fire,  a  part  of  the  carbon 
is  converted  into  carbon  monoxide,  which  escapes  unburned. 
This  results  in  the  loss  of  much  heat,  since  the  heat  of  combustion 
of  carbon  to  the  dioxide  is  97,000  calories  and  to  the  monoxide 
only  29,000  calories.  If  more  air  is  admitted  to  the  fire  than  is 
required  to  burn  the  coal,  heat  is  lost,  because  the  excess  air  carries 
away  heat  when  it  passes  up  the  chimney  with  the  products  of 


COAL,  COKE,  ILLUMINATING  GAS,  FLAMES  197 

combustion.  The  heat  lost  in  the  waste  gases  may  amount  to  as 
much  as  40  per  cent  of  the  calorific  value  of  the  coal  under  poor 
control  of  the  fire. 

When  large  quantities  of  coal  are  used  it  is  customary  to 
determine  the  efficiency  with  which  it  is  burned  by  analyzing  the 
flue  gases;  if  they  contain  carbon  monoxide  or  methane,  combus- 
tion has  been  incomplete  and  more  air  should  be  used ;  if  they  con- 
tain too  large  a  percentage  of  oxygen,  the  air  must  be  decreased. 
It  is  impossible  to  run  a  furnace  so  that  the  theoretical  amount  of 
air  is  used;  it  has  been  found  in  practice  that  about  twice  this 
amount  gives  the  best  results. 

223.  For  domestic  purposes  anthracite  coal  is  preferred  because 
it  can  be  handled  with  less  skill  and  attention.  It  is  hard  and  does 
not  crumble,  it  yields  less  volatile  matter  when  heated,  and  does 
not  produce  much  smoke;  and  since  it  does  not  melt  at  the  tempera- 
ture of  the  furnace  clinkers  do  not  form.  A  large  supply  of  coal 
can  be  placed  on  the  fire  without  smothering  it.  The  rate  at 
which  coal  burns  is  determined  by  the  amount  of  air  admitted  to 
it,  and  this  can  be  regulated  by  the  dampers  with  which  the  fur- 
nace is  supplied;  one  is  placed  below  the  grate  (A),  one  over  the 
fire  in  the  door  where  the  furnace  is  fed  (B),  and  one  in  the  pipe 
that  carries  off  the  products  of  combustion  (C).  When  a  fire  is 
started  a  strong  draft  is  desired  and  A  and  C  are  left  open  and  B 
closed;  by  this  arrangement  the  full  current  of  air  passes  through 
the  coal.  If  it  is  desired  to  check  the  fire,  A  is  placed  so  that  but  a 
small  amount  of  air  can  enter  the  furnace,  and  the  draft  produced 
by  the  hot  products  of  combustion  rising  in  the  chimney  is  reduced 
by  partially  closing  C0  If  an  insufficient  quantity  of  air  to  burn 
the  coal  is  admitted  to  the  furnace  in  this  way,  and  B  is  closed, 
carbon  monoxide,  methane,  and  hydrogen  sulphide  are  formed. 
The  air  burns  the  coal  resting  on  the  grate  to  carbon  dioxide,  and 
all  the  oxygen  is  consumed.  The  heat  generated  raises  the  next 
layer  of  coal  to  incandescence,  and  when  the  carbon  dioxide 
passes  over  it  carbon  monoxide  is  formed — C  +  CC>2  =  2CO. 
At  the  high  temperature  volatile  products  distill  from  the  coal  and 
since  no  oxygen  is  present  they  escape  unburned.  If  the  gases 
leak  through  a  crack  in  the  dome  of  the  furnace,  or  in  any  other 
way  get  into  the  current  of  hot  air  which  passes  around  the  fire- 
box and  is  led  through  pipes  to  the  rooms  heated  by  the  furnace, 


198  INORGANIC  CHEMISTRY  FOR  COLLEGES 

"  coal-gas  "  is  introduced  into  the  house.  This  soon  becomes  evi- 
dent as  the  result  of  the  action  of  the  hydrogen  sulphide  in  the  gas 
on  any  articles  made  of  silver  that  are  exposed  to  it;  they  tarnish 
as  the  result  of  the  formation  of  silver  sulphide.  If  the  furnace  is 
used  to  heat  water  or  steam  and  this  is  circulated  through  the  house, 
coal-gas  cannot,  of  course,  be  introduced  in  this  way.  But  even 
if  this  is  the  case  and  no  annoyance  is  caused  by  burning  the  coal 
incompletely,  the  method  is  uneconomical,  because  when  carbon  is 
burned  to  carbon  monoxide  only  about  one-third  of  the  heat  which 
can  be  produced  from  the  coal  is  obtained. 

When  coal  is  burning  in  an  insufficient  supply  of  air  and  the 
door  over  the  fire  is  opened,  there  is  usually  a  slight  puff  which  is 
caused  by  the  explosion  of  the  mixture  of  carbon  monoxide  and  air 
produced,  and  the  gas  then  burns  over  the  coal  with  a  blue  flame. 
The  evident  remedy  for  this  condition  is  to  admit  a  small  amount 
of  air  through  the  upper  door  of  the  furnace;  this  checks  the 
draught  through  the  fire,  and  thus  decreases  the  rate  at  which  the 
coal  burns,  and  it  furnishes  the  required  oxygen  to  burn  the 
carbon  monoxide  and  other  combustible  gases  produced. 

224.  Coke. — When  coal  is  heated  in  the  absence  of  air,  the  vola- 
tile materials  are  driven  off  and  the  residue  obtained,  consisting  of 
carbon  and  ash,  is  called  coke.  Large  quantities  of  coke  are  used 
in  extracting  iron  and  other  metals  from  their  ores.  A  small 
supply  is  obtained  as  a  by-product  in  making  illuminating  gas,  but 
the  large  amount  required  in  the  metallurgical  industries  is  pre- 
pared directly  for  this  purpose.  Coke  is  made  either  in  so-called 
bee-hive  ovens  (Fig.  20)  or  in  by-product  ovens.  The  former  are 
dome-shaped  ovens  of  brick,  about  10  feet  in  diameter  and  6  feet 
high,  provided  with  a  circular  opening  in  the  top  through  which  the 
coal  is  introduced  and  the  gaseous  products  formed  in  the  coking 
are  discharged.  There  is  also  an  opening  on  one  side  at  the  base  of 
the  oven,  which  serves  as  an  entrance  for  the  air  and  through 
which  the  coke  is  finally  withdrawn.  The  ovens  are  built  side 
by  side  and,  in  certain  districts  where  large  quantities  of  iron  are 
produced,  the  rows  of  ovens  are  a  quarter  of  a  mile  in  length.  An 
oven  is  charged  with  enough  coal  to  make  a  layer  2  feet  thick  and 
the  air-inlet  so  adjusted  that  just  enough  air  is  admitted  to  keep 
the  coal  red  hot.  The  heat  that  passes  through  the  walls  from  the 
two  ovens  on  either  side,  in  which  coke  is  being  made,  soon  starts 


COAL,  COKE,  ILLUMINATING  GAS,  FLAMES 


199 


the  combustion  of  the  coal.  By  charging  alternate  ovens  at  the 
proper  intervals  the  process  becomes  continuous.  It  takes  about 
two  days  to  complete  the  change  of  the  coal  to  coke.  The 
weight  of  the  latter  obtained  is  about  60  per  cent  of  the  weight  of 
the  coal.  The  process  is  very  wasteful  because  all  the  gaseous 
products  produced  escape  into  the  air  and  are  burned. 


FIG.  20. 

225.  Bee-hive  ovens  are  now  being  replaced,  in  part,  by  by- 
product ovens,  which  are  so  constructed  that  the  volatile  material 
formed  in  coking  is  saved.     When  this  material  is  cooled  it  yields 
combustible  gases,  and  a  mixture  of  liquids  and  solids  from  which 
benzene,  CeHe,  toluene,  CyHg,  and  many  other  valuable  substances 
are  obtained.     These  compounds  are  the  materials  from  which 
explosives,  dyes,  pharmaceutical  chemicals,  and  other  important 
products  are  made.     As  the  result  of  the  recent  war  the  produc- 
tion of   by-product  coke  was  markedly  increased   because  the 
by-products  were  needed  in  the  manufacture  of  explosives  and 
other  organic  compounds,  such  as  dyes,   which  had  been  pre- 
viously imported  from  Germany.     About  70  per  cent  of  the  coal 
used  is  recovered  as  coke  when  ovens  of  this  type  are  used.     About 
60  per  cent  of  the  gas  produced  is  used  to  heat  the  coal  and  the 
rest  is  utilized  as  a  source  of  power.     The  valuable  liquids  and 
solids  are  condensed,  separated,  and  purified. 

226.  A  by-product  oven  is  made  up  of  a  series  of  rectangular 
chambers  about  33  feet  long,  7.5  feet  high,  and  20  inches  wide, 


200 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


which  are  set  side  by  side  with  spaces  between  the  ovens  through 
which  the  hot  products  of  the  combustion  of  burning  gas  pass. 
The  volatile  products  escape  through  pipes  placed  at  the  top  of 
each  retort.  The  latter  has  a  door  at  either  end.  When  the  coal 
has  been  coked  these  are  opened,  and  the  coke  is  pushed  from  the 
furnace  by  a  mechanical  device.  The  products  of  a  by-product 
oven  are  about  70  per  cent  coke,  14  to  16  per  cent  gas  (about  9000 
cu.  ft.  per  ton),  4  to  6  per  cent  coal  tar,  which  contains  benzene, 
etc.,  and  0.25  to  0.30  per  cent  ammonia,  which  is  equivalent  to 
about  20  pounds  of  ammonium  sulphate  per  ton. 

227.  Illuminating  Gas. — Coal-gas  is  made  by  heating  bitumi- 
nous coal  in  retorts  made  of  fire-clay  about  8  feet  long,  18  inches 
wide,  and  15  inches  high.  Each  unit  consists  of  six  or  eight 

Hydraulic  Main 


fo-faru  Scrubber 
Condenser  i 


Gas  Holder. 


fteforf- 


FIG.  21. 

retorts  set  together  in  what  is  called  a  "  bench,"  and  each  is 
heated  either  by  burning  coke  on  a  grate,  or  by  gas.  The  gas 
produced  is  led  away  through  a  pipe  connected  with  each  retort ;  it 
is  first  allowed  to  cool  to  separate  out  the  tar,  and  is  then  passed 
through  "  scrubbers  "  containing  water,  which  dissolves  out  the 
ammonia.  It  is  then  brought  into  contact  with  slaked  lime  or  a 
hydrated  oxide  of  iron,  which  removes  sulphur  compounds  from 
the  gas,  and  is  finally  stored  in  a  gasometer.  (Fig.  21.)  The 
average  yields  of  the  products  obtained  from  1  ton  of  coal  are 
10,000  cu.  ft.  of  gas  of  16  candle  power,  1400  pounds  of  coke,  120 
pounds  of  tar,  and  5  pounds  of  ammonia. 

The  composition  of  the  gas  is  determined  by  the  temperature 
to  which  the  coal  is  heated,  1000°  to  1300°.  At  the  lower  tem- 
perature less  gas  is  obtained,  but  it  has  a  higher  candle  power. 
At  the  higher  temperature  the  gases  which  produce  light  when 


COAL,  COKE,  ILLUMINATING  GAS,  FLAMES  201 

they  burn  are  in  part  converted  into  others  which  give  little  light. 
The  former,  called  illuminants,  consist  principally  of  ethylene, 
C2H4;  acetylene,  C2H2;  and  benzene,  CeHe;  and  the  latter  are 
hydrogen  and  methane,  CH4.  Coal-gas  contains  ordinarily  about 
4  per  cent  illuminants,  49  per  cent  hydrogen,  35  per  cent  methane, 
7  per  cent  carbon  monoxide,  1  per  cent  carbon  dioxide,  and  4  per 
cent  nitrogen.  It  has  a  calorific  value  of  about  600  B.t.u.  per 
cubic  foot  and  the  specific  gravity  0.43  compared  with  air.  When 
a  gas  is  required  for  balloons  it  is  prepared  by  heating  coal  at  a 
very  high  temperature  so  that  the  chief  constituents  of  it  are 
hydrogen  and  methane. 

228.  Water-gas. — When  steam  is  passed  over  highly  heated 
carbon  a  reaction  takes  place  which  is  represented  by  the  following 
equation : 

C  +  H20  =  CO  +  H2  -  27,100  cal. 

This  reaction  is  utilized  in  making  water-gas.  Since  hydrogen 
burns  without  light  and  carbon  monoxide  with  a  blue  flame,  the 
gas  is  "  enriched  "  by  mixing  with  it  other  gases  obtained  by  heat- 
ing oils  to  a  high  temperature.  A  number  of  types  of  apparatus 
are  in  use  for  carrying  out  these  reactions.  In  one  of  these  anthra- 
cite coal  or  coke  is  brought  to  incandescence  in  a  generator  by 
means  of  a  blast  of  air.  (Fig.  22.)  The  hot  gases  produced  are  led 
through  the  carburetter,  which  is  built  of  fire-brick  and  is  filled  with 
a  "  checker-work  "  of  the  same  material  so  piled  that  the  bricks 
furnish  channels  for  the  gas  to  pass  through  and  present  a  large 
surface,  which  is  heated  white  hot.  The  gases  next  pass  through 
the  superheater,  similarly  constructed,  into  which  enough  air  is 
admitted  to  complete  the  burning  of  the  carbon  monoxide  formed 
in  the  generator.  The  checker-work  here  is  also  heated  to  incan- 
descence and  the  products  of  combustion  allowed  to  escape  from 
the  top.  When  this  condition  has  been  reached  the  air-blast  is 
cut  off  from  the  generator  and  superheated  steam  is  blown  through 
it.  The  mixture  of  carbon  monoxide  and  hydrogen  produced 
passes  through  the  carburetter,  to  which  is  admitted  a  slow  stream 
of  naphtha  or  other  oil.  This  is  vaporized  and  decomposed  and 
passes  along  with  the  gas  into  the  superheater  where  the  hydro- 
carbons are  converted  completely  into  gases  which  do  not  liquefy 


202 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


when  the  gas  cools.     The  enriched  or  carburetted  gas  which  issues 
from  the  superheater  is  scrubbed  and  stored. 

Since  the  water-gas  reaction,  C  +  H^O  =  CO  -f-  H2,  absorbs 
heat,  the  temperature  of  the  coke  in  the  generator  soon  falls  below 
that  at  which  the  reaction  takes  place,  which  is  about  1000.°  The 
steam  is  then  cut  off,  air  admitted  and  the  process  repeated.  Air 
is  usually  blown  for  about  eight  minutes  and  steam  for  six 
minutes,  the  gas  being  collected  only  when  the  steam  is  used. 
One  ton  of  anthracite  coal  yields  about  44,000  cubic  feet  of 
carburetted  gas,  which  contains  approximately  40  per  cent  of  hy- 
drogen, 17  per  cent  of  methane,  29  per  cent  of  carbon  monoxide, 


Coke 


'Smokestack 


Generator     Carbureter         Superheater 
FIG.  22. 

9  per  cent  of  illuminants,  4  per  cent  of  nitrogen,  and  1  per  cent 
of  carbon  dioxide.  The  gas  supplied  in  certain  cities  is  a  mixture 
of  coal-gas  and  water-gas. 

229.  Producer-Gas. — On  account  of  the  fact  that  gas  can  be 
used  conveniently  as  a  source  of  heat,  that  it  produces  no  ashes, 
and  is  very  efficient  as  a  source  of  power  when  used  in  an  explosive 
engine,  the  use  of  gas  on  the  large  scale  for  industrial  purposes  has 
greatly  increased  in  recent  years.  Since  only  the  calorific  value  of 
the  gas  is  of  importance  for  these  uses,  the  process  by  which  it  is 
made  is  simpler  than  that  employed  in  the  manufacture  of  illumi- 
nating gas.  In  the  simplest  kind  of  producer  to  make  this  kind  of 
gas  sufficient  air  is  blown  over  glowing  carbon  to  convert  it  into 
carbon  monoxide.  The  coal  used  is  supported  on  a  bed  of  ashes 


COAL,  COKE,  ILLUMINATING  GAS,  FLAMES  203 

resting  on  a  grate.  When  the  air  first  comes  in  contact  with  the 
burning  coal,  the  following  reaction  takes  place: 

C  +  O2  =  CO2  +  97,000  cal. 

The  heat  generated  is  carried  along  by  the  nitrogen  present  in  the 
air  and  by  the  carbon  dioxide  and  raises  the  coal  to  incandescence. 
This  then  reacts  with  the  carbon  dioxide  as  follows: 

CO2  +  C  =  2CO  -  39,000  cal. 

By  adding  these  two  equations  we  arrive  at  the  following,  which 
expresses  the  final  reaction : 

2C  +  O2  =  2CO  +  58,000  cal. 

A  large  amount  of  heat  is  generated  and  the  reaction  can  be  carried 
out  continuously.  The  carbon  monoxide  produced  can  be  used 
as  a. source  of  power  because  when  it  burns  heat  is  given  off: 

2CO  +  O2  =  2CO2  +  136,000  cal. 

Since  1  volume  of  oxygen  furnishes  2  volumes  of  carbon  monoxide — 
2C  +  O2  =  2CO — and  air  contains  4  volumes  of  nitrogen  to  1  of 
oxygen,  the  relation  between  the  volume  of  CO  and  N2  in  the  gas 
should  be  1  to  2.  That  this  theoretical  relation  can  be  closely 
reached  in  practice  is  seen  from  the  following  results  of  the  analysis 
of  a  producer-gas  made  as  outlined  above:  CO,  27.0;  N2,  55.3; 
C2H4,  2.5;  C2H2,  0.4;  H2,  12.0;  O2,  0.3;  and  CO2,  2.5  per  cent. 
The  hydrocarbons  and  hydrogen  are  produced  from  the  volatile 
matter  contained  in  the  coal  from  which  the  gas  was  prepared. 

Producer-gas  made  in  the  way  described  above  is  called  air-gas. 
When  large  quantities  of  gas  are  required  a  different  type  of  pro- 
ducer is  used  and  air  and  steam  are  blown  simultaneously  over  coal. 
The  heat  required  to  bring  about  the  water-gas  reaction  is  furnished 
by  that  produced  by  burning  the  carbon  to  carbon  monoxide. 
The  "  semi-water-gas  "  made  in  this  way  has  a  higher  calorific  value 
than  the  air-gas  because  it  contains  a  relatively  high  percentage 
of  hydrogen  and  much  less  nitrogen.  An  analysis  of  such  a  gas 
gave  the  following  results:  CO,  27.0;  N,  29.0;  CH4,  2.0;  H2, 
34;  CO2,  8.0.  Producer-gases  of  both  types  have  a  low  calorific 
value  on  account  of  the  large  percentage  of  nitrogen  which  they 


204 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


contain,  but  owing  to  their  ease  of  production  they  are  cheaper 
per  heat-unit  than  any  other  form  of  gas. 

230.  The  Burning  of  Gas. — The  chemical  changes  that  occur  in 
a  gas  flame  can  be  studied  conveniently  with  a  Bunsen  burner 
(Fig.  23) .  This  consists  of  a  tube  to  which  gas  is  admitted  at  one 
end  and  which  is  open  at  the  other,  where  the  issuing  gas  burns. 
Above  the  inlet  for  the  gas  there  are  two  holes  to  admit  air,  which 
can  be  opened  or  closed.  When  the  air  supply  is  closed,  ordinary 
illuminating  gas  burns  with  a  luminous  flame.  The  oxygen  of  the 
air  can  come  into  contact  with  the  issuing  gas  only  along  its  surface, 


-1610' 


FIG.  23. 


FIG.  24. 


and  as  combustion  takes  place  the  gas  within  the  cone  produced  is 
heated  to  a  high  temperature.  The  illuminants  present  in  the  gas, 
principally  ethylene  and  benzene,  are  decomposed  under  these  con- 
ditions into  hydrogen  and  carbon,  which  becomes  heated  to  such 
a  temperature  that  it  gives  off  light.  That  the  flame  contains 
carbon  can  be  shown  by  placing  a  cold  object  in  it;  carbon  is 
deposited  in  the  form  of  soot. 

When  the  holes  at  the  base  of  the  burner  are  opened,  air  is 
drawn  in  by  the  rising  gases  and  mixes  with  them.  The  flame  now 
consists  of  three  cones,  which  are  indicated  in  the  figure.  The 
lower  cone,  which  is  non-luminous,  consists  of  air  and  unburned  gas, 
and  as  the  mixture  rises  it  is  heated  by  the  burning  gas  which  sur- 


COAL,  COKE,  ILLUMINATING  GAS,  FLAMES  205 

rounds  it,  until  finally  ignition  takes  place.  The  boundary  between 
the  mixture  of  hot  gases  and  air  and  the  region  where  combustion 
begins,  is  the  surface  between  the  non-luminous  cone  and  second 
cone,  which  is  blue.  In  the  latter,  partial  combustion  of  the  gases  is 
taking  place.  This  has  been  determined  by  withdrawing  the  gases 
from  the  blue  cone  and  analyzing  them.  They  were  found  to 
consist  largely  of  carbon  monoxide  and  hydrogen  even  when  the 
gas  burned  was  pure  methane,  CEU:  2CH4  +  C>2  =  2CO  +  4H2. 
The  outer  zone  is  practically  non-luminous,  because  carbon  mon- 
oxide burns  with  a  light-blue  flame  and  hydrogen  burns  without 
the  evolution  of  light. 

The  fact  that  the  first  cone  contains  relatively  cool  gas  can  be 
shown  by  a  simple  experiment.  A  match  is  pierced  near  its  head 
by  a  pin  and  then  suspended  in  a  Bunsen  burner;  the  pin,  resting 
horizontally  across  the  top  of  the  burner,  supports  the  match, 
which  hangs  within  the  tube.  The  gas  is  now  lighted,  and,  if 
there  are  no  currents  of  air  to  blow  the  flame,  the  match  is  not 
ignited  for  some  time. 

The  temperatures  reached  in  different  parts  of  the  Bunsen 
flame  are  indicated  in  Fig.  23.  When  hydrogen  burns  in  pure 
oxygen  the  temperature  of  the  flame  is  about  2500°,  whereas  when 
it  burns  in  air  the  temperature  reached  is  much  lower.  If  equal 
weights  of  the  gas  are  burned  under  these  two  conditions,  the  same 
amount  of  heat  is  set  free  in  each  case;  when  oxygen  is  used,  the 
heat  generated  raises  the  temperature  of  the  water-vapor  pro- 
duced up  to  the  higher  temperature;  when  air  is  used,  the  large 
amount  of  nitrogen  present  is  also  heated  and,  as  a  consequence, 
the  temperature  does  not  rise  so  high.  It  is  evident  that  the 
highest  possible  temperature  attainable  when  hydrogen  or  any 
other  gas  is  burned  with  air  would  be  reached  when  the  latter 
contained  just  enough  oxygen  to  burn  the  gas.  It  is  impossible 
to  use  such  a  mixture  in  a  Bunsen  burner,  for  the  gas  would  not 
burn  quietly  at  the  mouth  of  the  burner;  the  mixture  would  ignite 
in  the  tube  and  the  gas  would  burn  where  it  enters  the  burner  below 
the  air  inlets.  If  too  much  air  is  admitted,  the  flame  "  strikes 
back  "  because  the  additional  amount  of  oxygen  produces  a 
mixture  in  the  tube  which  is  explosive  and  which  is  ignited  by 
the  gas  burning  at  the  mouth  of  the  burner.  It  is  for  this 
reason  that  only  a  part  of  the  oxygen  required  to  burn  the  gas 


206  INORGANIC  CHEMISTRY  FOR  COLLEGES 

is  admitted  at  the  base  of  the  burner  and  the  rest  is  obtained 
from  the  air  in  the  neighborhood  of  the  flame.  This  difficulty 
has  been  overcome  in  two  ingenious  ways.  In  the  blast-lamp  the 
gas  is  burned  from  a  tube  inside  of  which 'is  placed  a  second  tube 
through  which  air  is  forced  (Fig.  25).  The  gas  and  air  mix  at  the 
ends  of  the  tubes  and  there  is  no  opportunity  for  the  gas  to  strike 
back.  In  this  way  the  amount  of  air  necessary  to  burn  the  gas 
completely  can  be  mixed  with  it,  and  a  higher  temperature  can 
be  attained. 


FIG.  25. 

231.  A  comparatively  recent  invention  which  produces  the 
same  result  in  a  simpler  way  is  the  so-called  Meker  burner  (Fig.  24) . 
This  is  supplied  with  a  heavy  grid  of  metal  at  the  mouth  of  the 
burner  through  which  the  gases  pass.     When  properly  adjusted 
the  flame  cannot  strike  back,  because  the  gases  are  kept  below  their 
kindling-point  by  being  in  close  contact  with  the  metal  of  the  grid, 
which  is  comparatively  cold.     The  arrangement  is  a  novel  applica- 
tion of  the  principle  discovered  by  Davy  and  applied  by  him  to  the 
invention  of  the  miner's  safety  lamp.     A  much  larger  supply  of 
air  can  be  used  with  the  Meker  burner  than  with  the  Bunsen 
burner  and,  as  a  consequence,  the  amount  of  oxygen  mixed  with 
the  gas  is  nearer  that  required  for  complete  combustion  of  the 
gas  and  the  temperature  of  the  flame  is  higher. 

232.  Surface  Combustion. — The  rate  at  which  an  explosion 
travels  through  a  mixture  of  air  and  an    inflammable  gas  is  deter- 
mined by  the  proportions  of  the  two  present,  the  rate  being  the 
highest  when  there  is  just  enough  oxygen  to  burn  the  gas  com- 


COAL,  COKE,  ILLUMINATING  GAS,  FLAMES  207 

pletely.  It  is  possible  to  have  certain  mixtures  of  gases  move  so 
rapidly  that  their  rate  of  motion  is  greater  than  that  of  the  propa- 
gation of  the  explosion  through  them,  and  if  they  are  ignited  at  a 
distance  from  the  orifice  from  which  they  are  issuing,  they  will  burn 
and  the  flame  will  not  strike  back.  This  principle  is  used  in  sur- 
face combustion.  The  body  to  be  heated  is  surrounded  by  some 
refractory  material  in  granular  form  upon  the  surfaces  of  which  the 
gases  burn.  In  this  way  combustion  takes  place  where  the  heat  is 
desired,  a  higher  proportion  of  air  can  be  used  than  is  possible 
when  an  ordinary  burner  is  employed,  and  a  higher  temperature 
can  be  reached.  This  method  of  combustion  has  been  applied  to 
boilers  for  generating  steam,  by  filling  the  tubes  with  granular 
refractory  material  upon  the  surface  of  which  gas  is  burned. 

233.  Long-flame  Combustion. — When  an  inert  gas  like  nitro- 
gen or  carbon  dioxide  is  mixed  with  burning  gas  the  temperature 
of  the  flame  is  lowered,  as  we  have  seen.  The  size  of  the  flame  is 
increased  at  the  same  time,  for  owing  to  the  dilution  of  the  gases 
combustion  takes  place  more  slowly,  and  the  region  in  which  the 
burning  is  taking  place  is  larger.  This  interesting  scientific  fact 
has  been  utilized  only  recently;  its  application  to  industrial  uses 
has  resulted  in  marked  improvements  in  the  carrying  out  of  cer- 
tain manufacturing  operations.  One  example  is  of  particular 
interest.  Lime  is  made  by  heating  limestone — CaCOs  =  CaO  + 
CO2.  If  this  is  accomplished  by  means  of  burning  coal  the 
material  near  the  fire  is  heated  too  hot  and  that  at  some  distance 
from  it  is  not  completely  decomposed.  This  result  was  overcome  in 
the  old-style  kilns  by  mixing  fuel  with  the  limestone,  but  the  lime 
produced  contained  the  ashes  formed  and  was,  therefore,  impure. 
In  the  Eldred  kiln  a  part  of  the  carbon  dioxide  formed  from 
the  limestone  is  mixed  with  the  air  furnished  the  coal  on  the 
grate;  carbon  monoxide  is  formed  and  burns  when  mixed  with 
the  dioxide  in  a  large  flame  that  rises  through  the  kiln  and  pro- 
duces the  required  temperature  over  a  large  area.  So-called 
"  long-flame  "  combustion  is  a  marked  improvement  in  the  use  of 
gas  as  a  fuel  when  an  even  distribution  of  heat  is  required. 

EXERCISES 

1.  Two  samples  of  coal  were  analyzed  in  the  way  outlined  in  section  221. 
The  weights  given  below  were  obtained.     Calculate  in  each  case  the  per- 


208  INORGANIC  CHEMISTRY  FOR  COLLEGES 

centage  of  water,  volatile  matter,  fixed  carbon,  and  ash,  and  state  to  which 
class  each  coal  belonged.  In  the  case  of  one  sample  the  following  results  were 
obtained:  weight  of  crucible  which  contained  the  coal,  20.000  grams;  weight 
of  crucible  plus  coal,  21.000  grams;  weight  after  heating  at  100°,  20.986 
grams;  weight  after  driving  off  volatile  matter,  20.728  grams;  weight  after 
burning  fixed  carbon,  20.056  grams.  The  weights  obtained  with  the  second 
sample  were,  respectively  18.000  grams;  19.000  grams;  18.988  grams;  18.950 
grams;  18.069  grams. 

2.  Why  is  it  that  a  ton  of  bituminous  coal  gives  about  as  much  heat  as 
a  ton  of  anthracite  coal  or  coke,  although  it  contains  a  much  smaller  per- 
centage of  fixed  carbon? 

3.  (a)  Calculate  the  average  weight  of  the  molecules  in  a  mixture  con- 
taining oxygen  and  nitrogen  in  the  proportion  of  1  volume  of  the  former  to 
4  volumes  of  the  latter.     (6)  Calculate  the  average  weight  of  the  molecules 
in  coal-gas  assuming  the  average  molecular  weights  of  the  illuminants  to  be  40 
and  using  the  composition  of  coal-gas  given  in  section  227.     (c)  The  result 
obtained  in  a  is  approximately  the  average  weight  of  the  molecules  of  air. 
From  this  result  and  that  obtained  in  6  calculate  the  specific  gravity  of  coal- 
gas  compared  with  air  and  compare  the  result  with  the  figures  given  in  the 
text,     (d)  Using  the  composition  of  water-gas  given  in  section  228  calculate 
in  a  way  similar  to  that  just  used  the  specific  gravity  of  water  gas  compared 
with  air. 

4.  When  air-gas  is  made  in  a  producer  how  much  energy  is  lost  as  heat 
and  how  much  rendered  available  in  the  gas  produced?     Why  is  it  advantage- 
ous to  use  a  mixture  of  air  and  steam  in  the  producer?     In  this  case  how 
much  heat  is  lost  and  how  much  rendered  available  in  the  producer  gas? 

5.  (a)  If  a  balloon  having  the  weight  and  size  described  in  problem  12 
at  the  end  of  Chapter  V  were  filled  with  coal-gas,  what  weight  would  it  just 
lift  from  the  ground?     (6)  What  would  be  the  result  if  a  gas  composed  of 
equal  volumes  of  methane  and  hydrogen  were  used? 

6.  (a)  Calculate  the  number  of  calories  in  1  B.t.u.      (6)    Calculate  the 
factor  required  to  change  B.t.u.  per  pound  into  calories  per  gram. 

7.  Calculate  the  cost  of  1,000,000  B.t.u.  produced  by  burning    (a)  a 
sample  of  anthracite  coal  costing  $12  per  ton  and  furnishing  12,000  B.t.u. 
per  pound  and   (6.)  one  of  bituminous  coal,  costing  $8  per  ton  and  furnishing 
14,000  B.t.u.  per  pound, 


CHAPTER  XVII 

ACIDS,  BASES,  SALTS.     SOLUTIONS 

234.  As  the  chemical  behavior  of  the  substances  described  in 
the  preceding  chapters  has  been  discussed,  acids,  bases,  and  salts 
have  been  incidentally  mentioned.     It  is  advisable  at  this  point 
to  bring  together  the  isolated  facts  stated  here  and  there,  and  to 
discuss  more  fully  in  a  general  way  these  three  classes  of  com- 
pounds,  as    the   study    of    their   composition,    properties,    reac- 
tions, and  uses  constitutes  the  larger  part  of  inorganic  chemistry. 
The  attempt  to  interpret  the  behavior  of  these  substances  when 
dissolved  in  water  has  led  to  one  of  the  most  fruitful  theories  of 
modern  chemistry — a  theory  which  has  correlated  many  facts  of 
prime  importance  and  has  given  us  a  much  broader  and  deeper 
knowledge  of  chemical  phenomena. 

235.  Metallic  Elements. — A  more  or  less  definite  knowledge  of 
what  is  meant  by  the  word  metal  has  already  been  gained.     We 
associate  the  name  with  substances  which  possess  certain  physical 
properties  because  these  appeal  to  our  senses;    metals,  when  in 
compact  form,  have  a  surface  luster  which  is  so  characteristic  that 
it  is  defined  by  the  word  metallic;   they  are,  in  most  cases,  hard, 
ductile,  and  malleable,  and  are  good  conductors  of  heat  and  elec- 
tricity.    The  chemical  properties  of  metals  are  also  characteristic. 
They  form  oxides  and  hydroxides  which  dissolve  in  acids;  as  the 
result  of  the  reaction  salts  and  water  are  formed.     The  hydroxides 
of  the  metals  are  called  bases,  and  for  this  reason  metals  are  often 
called  base-forming  elements.     The  more  active  metals  react  with 
acids  and  hydrogen  is  set  free: 

Zn  +  2HC1  =  ZnCl2  +  H2 
Fe  +  H2SO4  =  FeS04  +  H2 

Both  the  physical  and  chemical  properties  of  metals  will  be  dis- 
cussed more  fully  later;  the  facts  stated  above  are  sufficient  for  the 
understanding  of  what  is  to  be  immediately  presented. 

209 


210  INORGANIC  CHEMISTRY  FOR  COLLEGES 

236.  Non-metallic  Elements. — The  elements  of  this  class  do  not 
possess  the  properties  characteristic  of  metals.     They  form  oxides 
which  react  with  water  to  produce  acids,  and  are,  accordingly, 
called  acid-forming  elements.     Their  hydroxides — the  compounds 
with   hydrogen   and   oxygen — are   acids.     Chlorine   and   sulphur 
are  typical  non-metallic  elements. 

237.  Bases. — The   hydroxides   of   the   metallic    elements  are 
bases;    those  which  dissolve  in  water,  such  as  sodium  hydroxide 
and  calcium  hydroxide,  are  called  alkalies.     Sodium  hydroxide 
and  potassium  hydroxide  are  caustic  alkalies,  and  they  are  so 
called  because  when  left  in  contact  with  the  skin  they  "  burn." 

Solutions  of  bases  in  water  affect  the  color  of  many  substances, 
and  those  which  change  readily  are  called  indicators.  The  indi- 
cators commonly  used  are  litmus,  which  is  changed  from  red  to 
blue  by  bases,  methyl-orange  which  changes  from  red  to  yellow, 
and  phenolphthalein,  a  substance  which  is  converted  into  a  red 
salt  by  bases. 

Solutions  of  bases  react  with  acids  and  form  salts,  and  as  the 
result  the  properties  of  the  base  disappear.  The  following  are 
equations  for  typical  reactions : 

NaOH  +  HC1  =  NaCl  +  H2O 

Ca(OH)2  +  H2SO4  =  CaSO4  +  2H2O 

238.  Acids. — All  acids  contain  at  least  one  hydrogen  atom  which 
can  be  replaced  by  the  more  active  metals.     They  have  a  sour 
taste,  affect  indicators,  and  react  with  metallic  oxides  and  hydrox- 
ides to  form  salts.     Acids  are  classed  as  monobasic,  dibasic,  and 
tribasic,  according  to  the  number  of  replaceable  hydrogen  atoms 
they  contain;  examples  are,  respectively,  HC1,  H2COs,  and  HaPC^. 
Acetic  acid,  C2H4O2,  contains  four  hydrogen  atoms,  but  it  is 
monobasic  because  of  these  one  only  can  be  replaced  by  metallic 
atoms;  for  this  reason  the  formula  is  sometimes  written  H  •  C2HsO2 
or  H(C2HsO2).     Acids  may  be  considered  as  made  up  of  replace- 
able hydrogen  atoms  and  acid  radicals,  which  are  groups  of  atoms 
that  remain  in  combination  when  the  acids  react  with  other  sub- 
stances.    The  radicals  of  carbonic  acid,  H2COs,  phosphoric  acid, 
H3PO4,  and  acetic  acid,  H-C2H3O2,  are  respectively,  COs,  PO4, 
and  C2H3O2. 


ACIDS,    BASES,    SALTS.    SOLUTIONS 


211 


239.  Salts.  —  When  an  acid  and  a  base  interact  to  form  a  salt 
the  reaction  is  said  to  be  one  of  neutralization;  the  resulting  sub- 
stance possesses  neither  the  characteristic  properties  of  the  acid 
nor  those  of  the  base.  When  a  solution  of  an  acid  is  neutralized 
by  one  of  a  base,  it  is  possible  to  tell  when  the  solution  is  neutral  if 
an  indicator  is  present;  at  first  the  color  produced  by  the  acid  is 
seen;  as  the  base  is  added  a  point  is  finally  reached  when  the  addi- 
tion of  a  single  drop  causes  a  per- 
manent change  in  color.  The  solu- 
tion is  neutral  when  a  drop  of  the 
acid  will  produce  one  color  of  the  <E= 
indicator  and  a  drop  of  the  base  will 
produce  the  other. 

240.  The  amount  of  acid  or  base  in 
a  solution  can  be  accurately  determined 
by  neutralizing  a  measured  volume  of  it 
with  the  aid  of  an  indicator.  Special 
forms  of  apparatus  have  been  devised 
for  this  purpose,  because  the  method  is 
much  used  in  analytical  chemistry.  The 
volumes  of  the  solutions  are  measured  in 
burettes,  which  are  long  tubes  of  glass 
provided  with  stop-cocks  at  one  end. 
They  are  marked  with  horizontal  lines, 
etched  in  the  glass,  so  that  the  volume 
of  the  solution  withdrawn  can  be  read 
(Fig.  26). 

An  example  will  be  instructive.     Sup- 
pose it  is  desired  to    determine  the  per- 
centage of  acid  in  a  sample  of  vinegar.    ^^^ 
Some  of  the  latter  is  put  into  a  burette,    w"/ 
and  a  definite  volume  is  drawn  off;  some 
water  and  an   indicator   are  now  added. 

A  solution  of  a  base,  the  strength  of  which  is  known,  is  next  run  in  cautiously 
from  a  burette  until  a  point  is  reached  where  one  drop  causes  a  permanent 
change  in  the  color  of  the  indicator.  The  volume  of  the  solution  of  the  base 
used  in  the  neutralization  is  noted,  and,  knowing  how  much  is  present  in 
1  c.c.,  the  weight  required  to  neutralize  the  known  volume  of  vinegar  can  be 
calculated.  By  using  the  chemical  equation  for  the  reaction  between  the 
base  and  the  acid  in  vinegar,  the  amount  of  acid  which  reacted  with  the  amount 
of  base  used,  can  be  determined,  and  from  this  the  percentage  of  the  acid  in  the 
sample  analyzed.  An  example  will  make  the  method  of  calculation  clear.  The 
volume  of  vinegar  taken  for  analysis  was  25  c.c.  To  neutralize  this  33.33  c.c. 
of  a  solution  of  sodium  hydroxide  which  contained  20  grams  of  the  base  in 


FIG.  26. 


212  INORGANIC  CHEMISTRY  FOR  COLLEGES 

a  liter  were  used.  Each  cubic  centimeter  of  the  base  contained  20  -J- 1000 
=  0.02  gram  of  NaOH.  Then  33.33  X  0.02  =  0.6666  gram  of  sodium 
hydroxide  was  used.  The  next  step  is  to  find  out  how  much  acetic  acid, 
which  is  the  acid  in  vinegar,  is  required  to  neutralize  this  weight  of  sodium 
hydroxide : 

0  6666  x 

NaOH        +          H-C2H302       =  NaC2H3O2  +  H2O 
23  +  16  +  1          1+24+3+32 
40  60 

40  :  60  :  0.6666  :x 
x  =  I 

The  acetic  acid  in  25  c.c.  of  vinegar  weighed,  accordingly,  1  gram.  Since  the 
specific  gravity  of  vinegar  is  approximately  equal  to  1,  the  percentage  of  acid 
in  the  sample  analyzed  was  1  -r-  25  =  0.04,  which  is  4  per  cent. 

241.  Salts  are  classed  as  neutral,  acid,  or  basic.     Neutral  salts 
are  formed  by  substituting  all  the  replaceable  hydrogen  of  the 
acid  by  metallic  atoms;  CaSCU,  A1PO4, and  Na2COa  are  formulas  of 
neutral  salts.     Since  some  so-called  neutral  salts  are  not  neutral 
to  indicators — a  fact  which  will  be  discussed  later — neutral  salts  are 
often  called  normal  salts. 

In  add  salts  but  a  part  of  the  hydrogen  in  the  acid  is  replaced 
by  metallic  atoms;  compounds  having  the  formulas  NaHSCU, 
NaHCOs,  and  Na2HP(>4  are  acid  salts.  A  number  of  methods  of 
naming  acid  salts  are  in  use.  Baking  soda,  for  example,  NaHCOs, 
may  be  called  acid  sodium  carbonate,  sodium  bicarbonate,  or 
sodium  hydrogen  carbonate.  When  an  acid  is  tribasic  it  may 
form  three  salts  with  one  base.  Phosphoric  acid  forms  salts  of 
the  composition  NaH2PO4,  Na2HPO4,  and  NasPCU;  the  first  is 
primary  sodium  phosphate  or  monosodium  phosphate,  the  second 
is  secondary  sodium  phosphate  or  disodium  phosphate,  and  the 
third  is  tertiary  sodium  phosphate  or  trisodium  phosphate. 

Basic  salts  usually  possess  compositions  which  are  more  or 
less  complicated.  A  neutral  salt  contains  a  metal  and  an  acid 
radical;  an  acid  salt  contains  hydrogen  in  addition  to  these;  a 
basic  salt  contains  a  metal,  an  acid  radical,  and  the  hydroxyl 
group,  the  group  characteristic  of  bases.  The  following  formulas 
represent  the  composition  of  some  basic  salts:  Fe(OH)SC>4, 
Pb3(OH)2(CO3)2,  and  Cu2(OH)2CO3. 

242.  Solutions. — It  has  already  been  indicated  that  acids  do  not 
react  with  metals  in  the  absence  of  water.     The  characteristic 


ACIDS,    BASES,    SALTS.    SOLUTIONS 


213 


properties  of  acids  due  to  the  hydrogen  atom  they  contain,  dis- 
appear in  the  absence  of  a  solvent.  Many  reactions  in  which 
bases  and  salts  take  part  are  dependent  on  the  presence  of  some 
liquid. 

Before  attempting  any  explanation  of  these  important  facts  it 
will  be  well  to  describe  some  striking  experiments  to  illustrate  the 
behavior  of  the  solutions  of  several  typical  substances  when  sub- 
mitted to  the  action  of  an  electric  current.  A  number  of  vessels 
are  provided  which  contain,  respectively,  water,  alcohol,  kerosene, 
pure  acetic  acid,  and  solutions  in  water  of  sugar,  hydrochloric  acid, 
sulphuric  acid,  sodium  hydroxide,  calcium  hydroxide,  sodium 


FIG.  27. 

chloride,  and  copper  sulphate.  An  apparatus  is  provided  to 
determine  whether  the  substances  contained  in  the  several  beakers 
will  allow  electricity  to  pass  through  them  (Fig.  27) .  This  consists 
of  an  electric  lamp  (a),  one  terminal  of  which  is  connected  with  a 
source  of  electricity  (e)  and  the  other  with  a  piece  of  platinum  foil  (6) 
to  serve  as  an  electrode.  The  latter  is  attached  to  a  non-conducting 
stand  to  which  is  joined  a  second  piece  of  platinum  (c)  to  serve 
as  the  other  electrode,  and  this  is  in  contact  with  a  wire  which 
returns  to  the  source  of  electricity. 

When  the  apparatus  is  connected  up  no  current  flows  on  ac- 
count of  the  break  in  the  circuit  at  the  electrodes.  If  a  vessel  (d) 
containing  a  solution  which  conducts  electricity  is  placed  so  that 
the  electrodes  dip  into  the  solution,  a  current  can  pass  and  the 
lamp  (a)  will  glow.  The  substances  enumerated  above  are  tested 
one  after  the  other,  the  electrodes  being  wiped  dry  after  each  test. 
It  will  be  found  that  there  is  no  evidence  that  either  water,  alcohol, 


214  INORGANIC  CHEMISTRY  FOR  COLLEGES 

kerosene,  or  pure  acetic  acid  conducts  the  current;  the  lamp  does 
not  glow.  Neither  does  the  presence  of  sugar  in  the  water  produce 
any  effect;  but  all  the  other  solutions  allow  the  current  to  pass 
and  the  lamp  glows  brightly.  The  substances  selected  for  the  ex- 
periment are  typical  and  represent  pure  compounds  and  solutions 
in  water.  The  conductivity  of  thousands  of  substances  has  been 
studied  and  the  results  are  in  accord  with  those  obtained  in  the 
experiment  just  described.  Solutions  of  acids,  bases,  and  salts 
conduct  the  electric  current  freely,  whereas  pure  compounds  or 
solutions  of  other  substances  such  as  sugar  do  not  conduct  at  all 
or  to  a  very  slight  degree.  Similar  results  are  obtained  if  certain 
solvents  other  than  water  are  used.  Although  solid  acids,  bases, 
and  salts  do  not  conduct  the  electric  current,  some  of  them  be- 
come conductors  when  they  are  in  the  molten  condition. 

Another  instructive  experiment  can  be  performed  to  emphasize 
the  effect  of  the  presence  of  water.  It  was  shown  that  the  lamp 
did  not  glow  when  the  electrodes  were  placed  in  water  or  in  pure 
acetic  acid,  which  is  a  colorless  liquid.  The  vessel  containing  the 
acid  is  placed  so  the  electrodes  dip  into  the  acid;  no  current  flows. 
Water  is  now  poured  into  it  slowly;  soon  the  lamp  begins  to  glow 
very  faintly ;  as  water  is  added  it  grows  brighter  and  brighter  and, 
finally,  almost  as  much  light  is  given  off  as  when  the  current  passed 
through  the  solution  of  sulphuric  acid.  These  results  are  of  great 
interest  and  need  an  explanation.  Water  does  not  conduct  the 
current  freely,  nor  does  acetic  acid;  but  the  mixture  does.  The 
solution  must  contain  something  which  is  not  present  in  either, 
before  they  are  mixed.  What  happens  when  an  acid,  salt,  or  base 
dissolves  in  water?  A  few  additional  facts  must  be  cited  before 
an  answer  is  given. 

When  an  electric  current  passes  through  a  solution  of  an  acid, 
base,  or  salt  a  profound  change  takes  place;  we  have  seen,  for 
example,  that  hydrochloric  acid  is  decomposed  into  hydrogen  and 
chlorine.  Electrolysis  takes  place,  and  we  say  that  the  current 
passes  as  the  result  of  electrolytic  conduction.  A  current  flows 
through  a  metal  without  producing  any  chemical  change;  in  this 
case  it  passes  as  the  result  of  what  is  called  metallic  conduction. 
If  electrolytic  conduction  in  a  solution  of  a  salt  is  studied,  it  will  be 
found  that  all  of  the  metal  can  be  set  free  at  one  pole  and  all  the 
acid  radical  at  the  other.  Take  the  case  of  copper  chloride.  At 


ACIDS,    BASES,    SALTS.    SOLUTIONS  215 

first  the  salt  is  uniformly  distributed  throughout  the  solution; 
at  the  end  of  the  electrolysis  all  the  copper  has  been  deposited  on 
one  pole  and  all  the  chlorine  set  free  at  the  other.  It  is  evident 
that  the  metallic  part  of  the  salt  is  attracted  by  one  electrode 
which  is  charged  negatively,  and  the  chlorine  by  the  other  which  is 
charged  positively.  It  is  the  same  with  other  salts;  the  metals 
and  the  acid  radicals  appear  to  be  independent  and  to  behave  dif- 
ferently when  in  solution;  and  they  are  attracted  by  the  elec- 
trodes which  are  charged  with  electricity. 

243.  The  Electrolytic  Dissociation  Theory. — Arrhenius,  a 
Swedish  chemist,  studied  critically  in  connection  with  the  proper- 
ties of  solutions  the  facts  noted  above  and  many  others,  and 
showed  that  they  could  be  interpreted  by  means  of  a  simple  hy- 
pothesis. It  was  necessary  to  explain  the  fact  that  solutions  of 
acids,  bases,  and  salts  conduct  the  electric  current,  that  in  these 
solutions  the  metals  and  hydrogen  behave  differently  from  the  acid 
radicals  and  the  hydroxyl  group,  but  that  they  all  are  attracted 
by  electrodes  charged  with  electricity.  The  assumption  made 
was  that  the  molecules  of  the  substances  which  dissolved — the 
solute — broke  down  into  smaller  particles,  called  ions,  and  that 
these,  as  a  result,  became  charged  with  electricity.  According  to 
this  view,  when  a  molecule  of  hydrogen  chloride  is  dissolved  in 
water,  the  latter  dissociates  the  molecule  into  a  hydrogen  atom 
which  becomes  positively  charged  with  electricity,  and  a  chlorine 
atom  which  becomes  negatively  charged.  This  is  represented  in 
symbols  thus: 

HCI  ->  H+  +  cr 

The  reasonableness  of  this  view  becomes  apparent  from  the  fol- 
lowing considerations.  When  any  two  different  substances  are 
brought  intimately  into  contact  and  are  then  separated,  they  each 
take  up  a  charge  of  electricity,  one  becoming  positively  and  the 
other  negatively  charged.  If  a  piece  of  silk  is  brought  into  close 
contact  with  a  rod  of  sulphur  by  rubbing,  and  the  two  are  sep- 
arated, both  are  found  to  be  charged  with  electricity.  An  experi- 
ment to  illustrate  this  fact  has  already  been  described  (7).  It  is 
not  unreasonable,  therefore,  that  in  some  similar  way  the  separa- 
tion of  two  atoms  may  produce  electricity  on  them.  The  water 
brings  about  the  separation,  and  as  it  offers  a  resistance  to  the 


216  INORGANIC  CHEMISTRY  FOR  COLLEGES 

tendency  of  positive  and  negative  electricity  to  unite,  the  charged 
atoms  remain  apart.  In  general,  the  greater  the  resistance 
shown  by  a  liquid  to  this  tendency,  the  more  it  dissociates  sub- 
stances. The  atoms  charged  with  electricity  which  are  formed 
in  this  way  are  called  ions. 

The  ion  of  chlorine  is  assumed  to  be  an  atom  of  chlorine  com- 
bined with  a  certain  amount  of  electricity;  it  differs  in  energy 
from  a  simple  chlorine  atom  and  should  have  different  properties 
from  those  of  the  latter.  We  have  seen  that  the  amounts  of 
energy  associated  with  carbon  in  its  two  forms,  diamond  and 
graphite,  are  different.  We  are,  accordingly,  not  surprised  to 
find  that  a  chlorine  ion  possesses  different  properties  from  chlorine 
gas,  and  it  is  entirely  reasonable  that  when  electricity  is  removed 
from  the  atom  and  its  energy  thereby  changed,  the  element  should 
exhibit  the  properties  with  which  we  are  familiar.  The  hypothe- 
sis put  forward  by  Arrhenius  has  served  to  explain  so  many  appar- 
ently unrelated  facts  that  it  is  now  one  of  the  most  fruitful  theories 
of  modern  chemistry.  It  is  called  the  electrolytic  dissociation 
theory,  because,  as  the  name  implies,  it  involves  the  assumption  of 
a  dissociation  of  molecules  into  particles  charged  with  electricity. 

Sugar  when  dissolved  does  not  conduct  the  electric  current; 
accordingly,  it  is  called  a  non-electrolyte,  and  does  not  dissociate 
into  ions.  Acids,  bases,  and  salts  are  called  electrolytes  because 
their  solutions  in  water  conduct  an  electric  current.  The  ions  of 
monobasic  acids  are  hydrogen  and  the  acid  radical,  thus  the  ions 
of  HC1  are  H+  and  Cl~,  of  nitric  acid  H+  and  NOs".  The  ioniza- 
tion  of  dibasic  and  tribasic  acids  takes  place  in  steps.  Sulphuric 
acid  dissociates  as  follows: 


+H+  +  HSO4~ 
HS04--»H+  +  S04~ 

The  complete  ionization  of  the  acid  is  represented  thus: 
H2SO4.->2H+  +  SO4~ 

It  should  be  noticed  that  each  hydrogen  atom  carries  one  positive 
charge  and  the  number  of  charges  on  any  atom  or  group  is  the 
same  as  its  valence.  The  PO4  ion  from  phosphoric  acid,  HaPC^,  is 


ACIDS,    BASES,    SALTS.    SOLUTIONS  217 

represented  thus,  PO4         .    When  any  substance  undergoes  ioniza- 
tion  the  number  of  positive  charges  always  equals  the  number  of 
negative  charges — a  fact  which  should  be  carefully  noted  when 
ionic  formulas  are  used  in  writing  equations. 
Bases  produce  metallic  ions  and  hydroxyl  ions : 

NaOH  -»  Na+  +  OH~,     and     Ca(OH)2  ->  Ca+  +  +  20HT 

The  ions  of  salts  are  formed  from  the  metallic  atoms  and  the  acid 
radicals : 

NaCl->Na+    +  Cl~ 

Na2SO4  ~*  2Na+  +  SO*" 
CaSO4-*Ca+  +  +  SO*" 

It  is  noted  again  that  the  number  of  charges  equals  the  valence  of 
the  atom  or  group;  calcium  is  bivalent  and  its  ion  is  accordingly 
Ca++.  Metals  and  hydrogen  form  positive  ions,  and  acid  radicals 
and  the  hydroxyl  group  form  negative  ions.  For  this  reason 
metals  are  sometimes  spoken  of  as  positive  elements  and  non- 
metals  as  negative  elements. 

From  the  standpoint  of  the  electrolytic  dissociation  theory 
the  characteristic  properties  of  an  acid  are  due  to  the  hydrogen 
ion,  for  these  properties  are  evident  only  when  the  acid  is  in  solu- 
tion and  when  ions  are  formed.  According  to  this  view  an  acid  is  a 
substance  which  yields  one  or  more  hydrogen  ions  when  it  is  dis- 
solved in  water.  The  characteristic  properties  of  bases  are  due  to 
the  hydroxyl  ion,  OH~;  consequently,  a  base  is  described  as  a 
substance  which  yields  a  hydroxyl  ion  when  dissolved  in  water. 

244.  Interpretation  of  Typical  Reactions  with  the  Use  of  the 
Theory  of  Ions.  Electrolysis. — When  hydrochloric  acid  is  dis- 
solved in  water  it  is  assumed  that  it  breaks  down  into  hydrogen 
ions,  which  are  positively  charged  atoms,  and  chlorine  ions,  which 
are  negatively  charged  atoms.  The  terminals  of  a  source  of  elec- 
tricity, a  storage  battery,  for  example,  are  charged,  one  positively 
and  one  negatively.  When  these  are  connected  with  the  elec- 
trodes (a)  and  (6)  (Fig.  28)  which  dip  into  a  solution  of  hydro- 
chloric acid,  we  have  in  contact  with  the  solution  two  poles  that  are 
charged  with  electricity.  The  positive  charge  on  a  attracts  the 


218 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


^:~^— 


negatively  charged  chlorine  ions  indicated  by  ©  in  the  diagram. 
As  each  ion  comes  in  contact  with  the  electrode  it  is  discharged, 
for  when  positive  and  negative  electricity  come  together,  the  elec- 
tricity disappears.  As  the  result,  in  this  case,  the  chlorine  ion 
loses  its  charge  and  changes  to  an  atom  from  which  molecules  of 
chlorine  are  formed;  and  the  gas  as  we  know  it  is  set  free.  At  the 
same  time  positive  electricity  flows  from  the  battery  to  take  the 
place  of  that  which  disappeared  when  it  united  with  the  negative 
charge  on  the  chlorine  ion.  As  the  result  of  the  continual  discharge 

of  the  electrode  in  this  way,  chlorine 
gas  is  formed,  and  the  battery  furnishes 
electricity  to  the  pole;  as  a  conse- 
quence, a  current  passes  through  the 
wire  connecting  the  battery  with  the 
electrode  as  indicated  by  the  arrow. 

Similar  changes  take  place  at  the 
other  pole.  The  positively  charged 
hydrogen  atoms  give  up  their  charge 
to  the  negative  electrode;  hydrogen  is 
set  free,  and  the  battery  keeps  the 
electrode  negatively  charged. 

As  the  electrolysis  proceeds  the 
negative  ions  move  to  the  positive 
pole,  which  is  called  the  anode  and  the 
positive  ions  to  the  negative  pole,  the 
cathode,  and  finally  all  the  ions  are 
discharged  and  the  current  ceases  to 

flow.  The  changes  described  have  been  studied  quantitatively 
and  generalizations  of  the  greatest  importance  have  been  dis- 
covered; this  will  be  considered  at  length  later. 

When  an  electric  current  is  passed  through  a  solution  of 
sulphuric  acid,  hydrogen  and  oxygen  are  obtained  (41).  In  this 
case  a  secondary  reaction  takes  place  at  the  anode.  The  negative 
ion,  SC>4~~,  is  discharged,  but  unlike  chlorine  it  does  not  exist  as 
a  chemical  compound;  it  immediately  reacts  with  the  water, 
2SO4  +  2H2O  =  2H2SO4  +  O2,  and  oxygen  is  set  free.  A  similar 
reaction  occurs  in  most  cases  when  an  acid  radical  which  cannot 
exist  hi  the  free  condition  is  liberated ;  it  changes  into  the  acid  by 
uniting  with  hydrogen,  and  oxygen  is  set  free.  The  acid  is  thus 


FIG,  28. 


ACIDS,    BASES,    SALTS.    SOLUTIONS  219 

continually  regenerated.  It  was  stated  that  in  the  decomposition 
of  water  into  hydrogen  and  oxygen,  sulphuric  acid  acted  as  a 
catalytic  agent  (27);  the  manner  in  which  it  acts  is  now  clear. 
The  hydrogen  and  oxygen  come  from  the  water  but  the  latter  first 
reacts  with  the  discharged  SO4  ion  and  forms  sulphuric  acid, 
which  is  decomposed  by  the  current. 

245.  Neutralization.  The  following  equations,  in  which  the 
ions  involved  are  represented,  bring  out  relationships  in  neutraliza- 
tion which  would  not  be  apparent  without  the  theory  of  electro- 
lytic dissociation. 

Na+  +  OH-  +  H+  +  OP  =  Na+  +  CP  +  H2O 

K+  +  OH-  +  H+  +  cr  =  K+  +  cr  +  H2o 

Ca+  +  +  20H~  +  2H+  +  2C1~  =  Ca++  +  2C1~  +  2H2O 
Ca++  +  20H-  +  2H+  +  SO4~"~  =  Ca++  +  SO4~"  +  2H2O 

In  each  case  the  reaction  consists  in  the  union  of'  hydrogen  ions 
with  hydroxyl  ions  to  form  molecules  of  water,  which  are  undis- 
sociated.  A  sodium  ion,  Na+,  is  represented  on  the  left-hand  side 
of  the  first  equation;  it  was  formed  when  sodium  hydroxide  was 
dissolved  in  water.  A  sodium  ion  is  represented  on  the  right-hand 
side  of  this  equation.  From  the  viewpoint  of  ions  the  sodium  has 
not  taken  part  in  the  reaction  at  all;  likewise  the  chlorine  ion  has 
been  inactive.  We  do  not  write  NaCl  on  the  right  side  of  the 
equation  to  indicate  that  sodium  chloride  is  formed ;  in  solution  the 
salt  is  ionized  and  should  appear  as  Na+  +  Cl~. 

Whether  such  a  view  is  correct  can  be  tested  experimentally. 
In  the  reactions  represented  by  the  first  two  equations  above  the 
only  chemical  change  is  the  formation  of  water  from  hydrogen 
and  hydroxyl  ions;  if  this  is  true  the  heat  produced  when  the  quan- 
tities represented  in  the  equation  are  brought  together  should  be 
the  same.  If  1  gram-molecule  of  sodium  hydroxide  (23  +  16  +  1 
=  40  grams)  is  dissolved  in  water  and  neutralized  with  a  solution  of 
hydrochloric  acid  (1  +  35.5  =  36  grams),  13,600  calories  are 
evolved.  If  a  similar  experiment  is  carried  out  and  1  gram-mole- 
cule of  potassium  hydroxide  (39  +  16+1  =56  grams)  is  used, 
the  same  amount  of  heat  is  set  free.  When  the  heat  evolved  in  the 
reactions  represented  by  the  third  and  fourth  equations  above  is 


220  INORGANIC  CHEMISTRY  FOR  COLLEGES 

determined  in  the  same  way,  it  is  found  that  2  X  13,600  calories 
are  set  free  in  each  case,  for  2  gram-molecules  of  water  are  formed. 
The  heat  of  neutralization  is  a  constant  and  can  be  represented 
thus: 

H+  +  OH~  =  EkO  +  13,600  cal. 

This  statement  holds  true  only  when  solutions  containing  com- 
pletely dissociated  acids  and  bases  are  used.  We  shall  see  later 
that  under  certain  circumstances  acids,  bases,  and  salts  are  only 
partially  broken  down  into  ions.  The  fact  that  the  heat  of  neu- 
tralization is  a  constant  is  strong  evidence  of  the  correctness  of 
the  conception  of  ions. 

246.  Acids  ,  bases,  and  salts  have  each  two  distinct  sets  of  prop- 
erties. We  have  seen  that  certain  properties  of  acids  can  be 
attributed  to  the  hydrogen  ions  which  they  yield  when  dissolved 
in  water.  Each  acid  also  possesses  properties  which  are  deter- 
mined by  the  acid  radical  it  contains;  thus  a  solution  of  hydro- 
chloric acid  contains  hydrogen  ions,  H+,  and  chlorine  ions,  Cl~~, 
and  the  properties  of  the  latter  are  shown  by  the  acid.  Likewise 
sodium  chloride  gives  a  sodium  ion  and  a  chlorine  ion,  when  dis- 
solved in  water.  If  this  view  is  correct  its  solution  should  possess 
the  same  set  of  properties  which  hydrochloric  acid  possesses  due 
to  the  fact  that  it,  also,  yields  a  chlorine  ion.  Or  making  the 
conclusion  more  general,  all  substances,  whatever  they  may  be, 
should  show  certain  properties  in  common  provided  they  all  yield  a 
chlorine  ion  when  dissolved  in  water.  Such  a  conclusion  is  a  logical 
consequence  of  the  theory  —  and  the  facts  are  in  accord  with  the 
theory.  We  have  learned  that  a  test  that  can  be  applied  to  all  the 
chlorides  of  the  metals  is  to  treat  their  solutions  with  a  solution  of 
silver  nitrate;  in  all  cases  silver  chloride  is  formed  —  they  all  show 
one  property  in  common.  Ionic  equations  for  the  reactions  serve 
to  emphasize  this  fact  : 

K+  +  Cr  +  Ag+  +  NO3~   =  AgCl  +  K+  +  N03~ 


Ag+  +  NO3~  =  AgCl  +  Na+  +  NO 


The  reaction  in  the  two  cases  represented  above  consists  in  the 
union  of  the  silver  and  chlorine  ions  to  form  silver  chloride,  which 
separates  as  a  precipitate;  this  occurs  as  the  amount  of  the  chlo- 


ACIDS,    BASES,    SALTS.    SOLUTIONS  221 

ride  that  can  be  formed  ordinarily  when  solutions  are  brought 
together  is  greater  than  the  amount  which  will  remain  dissolved  in 
water  (1  liter  of  water  dissolves  0.0013  gram  AgCl).  Silver  chlo- 
ride is  written  in  the  non-ionized  condition  as  AgCl,  because  it 
separates  as  a  precipitate;  ions  exist  in  solutions  only,  therefore  in 
writing  ionic  equations,  substances  not  dissolved  are  expressed 
by  their  molecular  formulas.  From  the  above  it  is  seen  that  the 
test  for  chlorides  (145)  is  a  test  for  chlorine  ions;  it  can  be  made 
not  only  with  silver  nitrate,  but  any  soluble  silver  compound  which 
gives  a  silver  ion. 

Like  acids,  salts  and  bases  when  in  solution  have  two  distinct 
sets  of  properties;  all  sodium  salts,  for  example,  have  a  set  of 
properties  in  common  upon  which  a  test  for  sodium  can  be  based. 
The  facts  enumerated  above  find  a  satisfactory  explanation  in  the 
theory  of  ions;  they  were  known  before  the  theory  was  put  for- 
ward, but  no  reasonable  interpretation  of  them  was  given. 

247.  Reactions  of  Double  Decomposition.  —  The  application  of 
the  theory  of  electrolytic  dissociation  to  reactions  of  neutralization 
and  to  the  test  for  chlorides  can  be  broadened  to  include  all  other 
cases  of  double  decomposition.  A  few  equations  written  with  ionic 
symbols  will  make  this  clear: 


2Na+  +  CO3~     +  Ca++  +  2C1"   =  2Na^  +  2C1"  +  CaCO3 
Fe+  +  +  2Cr  +  2Na+  +  2OH~  =  2Na+  +  2C1~  +  Fe(OH)2 
2K+  +  20H~  +  2H+  +  SO4—  =  2K+  +  SO4"     +  2H2O 
2Na+  +  2Cr  +  2H++  SO4~~  =  2Na  +  +  SO4~"  +  2HC1 

In  all  cases  of  double  decomposition  at  least  one  substance  escapes 
from  the  system  undergoing  change.  In  the  equations  given  above 
CaCOs  and  Fe(OH)2  are  written  in  the  molecular  form  because 
they  are  insoluble  substances  and  precipitate;  they  are  not  in 
solution  and,  therefore,  not  in  the  ionic  condition.  The  formula 
of  hydrochloric  acid  is  written  in  the  molecular  form  because  it 
escapes  as  a  gas.  The  formula  for  water  is  also  molecular  because 
it  does  not  appreciably  dissociate.  It  has  been  stated  (149)  that  a 
double  decomposition  takes  place  if  one  of  the  products  is  a  gas  or  is 
insoluble.  A  third  condition  can  now  be  added;  reactions  of  this 
type  take  place  if  one  of  the  products  is  a  substance  which  is  un- 


222  INORGANIC  CHEMISTRY  FOR  COLLEGES 

dissociated  or  slightly  dissociated.  It  is  seen  from  the  above  that 
if  we  wish  to  make  any  insoluble  acid,  base,  or  salt  we  can  do  so  by 
bringing  together  in  solution  two  substances  one  of  which  yields  the 
positive  and  the  other  the  negative  ion  of  the  compound  sought. 

248.  The  Reaction  between  Metals  and  Acids. — The  reaction 
between  zinc  and  sulphuric  acid  is  represented  with  ionic  symbols 
as  follows: 

Zn  +  2H+  +  S04~"  =  Zn^  +  H2  +  S04~ 

Hydrogen  gas,  H2,  and  zinc  sulphate,  ZnSC>4,  are  formed;  as  the 
salt  is  dissolved  in  the  water  used  as  a  solvent  it  is  represented  as 
dissociated  into  ions.  The  only  change  which  takes  place  is, 
according  to  this  view,the  transfer  of  the  electric  charge  from  the 
hydrogen  to  the  zinc.  When  copper  is  treated  with  a  solution  of 
sulphuric  acid  a  similar  change  does  not  take  place;  this  is  explained 
by  saying  that  hydrogen  has  a  greater  tendency  to  be  an  ion  than 
copper,  and  does  not  give  up  its  charge  to  the  metal.  With  zinc, 
however,  the  case  is  different;  it  has  a  greater  tendency  to  be  an 
ion  than  hydrogen  and,  accordingly,  passes  into  solution  and  forces 
the  hydrogen  out. 

All  elements  tend  more  or  less  to  pass  from  the  condition  of 
the  free  element  into  that  of  an  ion;  they  differ  markedly  among 
themselves  in  the  strength  of  this  tendency,  which  determines  the 
chemical  behavior  of  the  elements  in  many  of  their  reactions  with 
other  substances. 

249.  The  Reaction  between  Metals  and  Salts. — In  the  cases  cited 
the  tendency  of  zinc  to  form  ions  is  greater  than  that  of  hydrogen, 
and  that  of  copper  is  less  than  that  of  hydrogen.     We  might 
well  ask  what  would  happen  if  metallic  zinc  were  put  into  a  solution 
of  copper  sulphate.     The  experiment  can  be  readily  performed, 
and  the  results  are  what  is  anticipated ;  they  are  represented  by 
the  following  equation : 

Zn  +  Cu++  +  S04--  =  Zn++  +  Cu  +  SO4— 

The  greater  tendency  of  zinc  to  form  ions  asserts  itself;  it  passes 
into  solution  and  metallic  copper  is  precipitated. 

250.  The  Reaction  between  Hydrogen  and  Salts. — Another  case 
of  replacement  of  this  kind  is  of  interest.     Platinum,  we  have 
learned,  is  a  very  inactive  metal;  it  forms  ions  less  readily  than 


ACIDS,  BASES,  SALTS.    SOLUTIONS  223 

hydrogen  and,  consequently,  does  not  replace  it  when  the  metal 
is  brought  into  contact  with  an  acid.  We  might  prophesy  what 
would  happen  if  hydrogen  gas  were  passed  through  a  solution  of 
platinum  chloride,  in  which  platinum  ions  and  chlorine  ions  are 
present.  The  greater  ionizing  tendency  of  hydrogen  comes  into 
play;  it  becomes  the  ion  and  the  platinum  is  forced  out  of  solu- 
tion. The  equation  for  the  reaction  is  as  follows: 

Pt++  +  2C1"  +  H2  =  Pt  +  2Cr  +  2H+ 

When  hydrogen  gas  is  bubbled  through  a  solution  of  platinous 
chloride,  platinum  is  precipitated. 

251.  The  Reaction  between  Non-metals  and  Salts. — One  more 
case  of  replacement,  in  which  negative  ions  are  involved,  will  serve 
to  emphasize  the  principle  under  discussion.     It  will  be  recalled 
that  in  one  of  the  tests  for  free  chlorine  the  gas  is  treated  with  a 
solution  of  potassium  iodide  (130);  iodine  is  set  free  and  is  recog- 
nized by  its  action  on  starch.    According  to  the  conception  of  ions 
the  reaction  takes  place  as  indicated  by  the  following  equation: 

C12  +  2K+  +  2I~  =  2Cr  +  2K+  +  I2 

When  chlorine  passes  into  solution  it  becomes  a  negative  ion,  since 
it  is  a  non-metallic  element.  Its  tendency  to  pass  into  the  ionic 
condition  is  greater  than  that  of  iodine,  and  when  brought  into 
contact  with  iodine  ions  the  exchange  in  the  electric  charge 
takes  place,  and  iodine  is  set  free  in  the  elementary  condition. 
Elaborate  investigations  have  been  made  of  the  replacing  power 
of  elements,  because  a  measure  of  the  tendency  of  an  element  to 
pass  into  the  ionic  condition  is  an  index  not  only  of  its  ability  to 
replace  other  elements,  but  of  its  general  chemical  activity.  The 
elements  which  form  ions  readily  are  the  ones  we  have  called 
active  elements,  because  a  large  amount  of  energy  is  set  free 
when  they  react  with  other  substances. 

252.  Electromotive  Series  of  the  Metals. — There  are  a  number 
of  independent  ways  of  measuring  the  tendency  of  a  free  element 
to  pass  into  solution  as  an  ion,  and  they  all  give  approximately  the 
same  result.     One  method  is  to  place  samples  of  a  given  metal 
into  solutions  of  salts  of  other  metals  and  determine  if  a  reaction 
takes  place.     For  example,  if  some  metallic  mercury  is  placed  in  a 
solution  of  silver  nitrate  it  will  be  found  that  the  former  passes 


224 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


into  solution  and  silver  separates  out ;  the  tendency  of  mercury  to 
to  form  ions  is,  therefore,  greater  than  that  of  silver.  By  a  series 
of  experiments  it  is  possible  to  find  out  which  metals  possess  this 
property  to  a  greater  or  less  degree  than  any  given  metal.  By 
carrying  out  experiments  in  this  way  with  each  metal  it  is  pos- 
sible to  make  a  table  in  which  the  elements  are  arranged  according 
to  their  relative  tendencies  to  form  ions.  There  is  given  for 
reference  on  this  page  such  a  table,  in  which  the  arrangement 
is  according  to  the  decreasing  values  for  the  com- 
mon metals.  It  will  be  seen  that  what  we  have 
called  the  active  metals  are  at  the  top  of  the  list. 
All  the  metals  above  hydrogen  set  it  free  from 
acids,  whereas  those  below  hydrogen  do  not.  We 
shall  see  as  the  chemistry  of  the  metals  is  studied 
that  the  position  of  an  element  in  the  series  is,  in 
general,  an  indication  of  its  chemical  behavior. 
This  arrangement  of  the  metals  is  called  the  elec- 
tromotive series  because  it  was  arrived  at  as  the 
result  of  the  study  of  certain  electrical  properties 
of  solutions  which  will  be  discussed  later. 

253.  The  Boiling-point  of  Solutions. — Pure  water 
boils  at  100°  at  760  mm.  If  a  solid  is  dissolved  in  it, 
it  will  be  found  that  the  solution  must  be  heated  to 
a  higher  temperature  before  it  boils.  For  example, 
the  boiling-point  of  water  can  be  raised  to  108° 
by  dissolving  salt  in  it.  The  surface  of  the  solution 
contains  a  smaller  number  of  molecules  of  water 
than  an  equal  surface  of  pure  water.  In  order  to 
have  the  same  number  of  molecules  leave  the  two 
surfaces  in  the  same  time  and  thus  produce  the 
same  vapor-pressure  (179),  it  is  necessary  to  raise 
the  temperature  of  the  solution  to  impart  a  greater 
velocity  to  the  molecules  of  water  in  the  solution. 

The  effect  of  dissolved  substances  on  the  boiling-points  of 
liquids  was  carefully  investigated  by  Raoult,  a  French  scientist,  who 
discovered  facts  of  the  greatest  interest  and  importance.  He 
found  that  there  appeared  to  be  no  regularity  in  the  behavior  of 
solutions  of  acids,  bases,  and  salts,  but  that  the  effect  on  the  boiling- 
point  of  what  we  now  call  non-electrolytes  was  such  that  it  could  be 


Potassium 

Sodium 

Barium 

Strontium 

Calcium 

Magnesium 

Aluminium 

Manganese 

Zinc 

Chromium 

Cadmium 

Iron 

Cobalt 

Nickel 

Tin 

Lead 

Hydrogen 

Copper +  + 

Arsenic 

Bismuth 

Antimony 

Mercury + 

Silver 

Palladium 

Platinum 

Gold 


ACIDS,    BASES,    SALTS.    SOLUTIONS  225 

expressed  by  two  simple  laws.  The  first  of  these  states  that  the 
rise  in  boiling-point  of  a  solution  of  a  given  non-electrolyte  is  pro- 
portional to  the  weight  of  the  dissolved  substance  provided  the 
same  weight  of  solvent  is  taken  in  each  case.  For  example,  when 
100  grams  of  glycerine  are  dissolved  in  1  liter  of  water  the  boiling- 
point  of  the  solution  is  100.56°;  if  50  grams  are  dissolved  in  the 
same  amount  of  water  the  boiling-point  is  100.28°,  the  rise  being 
one-half  that  produced  by  100  grams.  The  second  law  is  as  fol- 
lows: The  boiling-points  of  all  solutions  of  non-electrolytes  in  the 
same  solvent  are  the  same,  provided  the  amounts  dissolved  in  a 
given  weight  of  the  solvent  are  in  the  ratio  of  the  molecular  weights 
of  che  solutes.  The  molecular  weights  of  sugar,  Ci2H22On, 
glycerine,  CaHgOs,  and  glucose,  C6Hi2O6,  are  respectively  343, 
92,  and  180.  If  343  grams  of  sugar  are  dissolved  in  1000  c.c.  of 
water  the  solution  will  boil  at  100.52°;  if  92  grams  of  glycerine  or 
180  grams  of  glucose  are  dissolved  in  1000  c.c.  of  water  the  boiling- 
point  in  either  case  will  also  be  100.52°. 

Neither  of  the  above  laws  holds  for  electrolytes;  for  example, 
a  solution  which  contains  1  gram-molecule  of  salt,  NaCl  (58.5 
grams),  dissolved  in  1000  c.c.  boils  at  100.87°.  No  explanation  of 
"  the  fact  that  the  behavior  of  solutions  of  acids,  bases,  and  salts 
is  not  in  accord  with  these  laws  was  furnished  until  the  electrolytic 
dissociation  theory  was  put  forward.  We  shall  see  that  this  theory 
supplied  not  only  a  reasonable  interpretation  of  the  observed 
results,  but  opened  up  the  way  for  a  detailed  study  of  solutions  from 
a  new  point  of  view,  which  has  increased  markedly  our  knowledge 
of  chemistry. 

If  the  weight  of  a  molecule  of  a  certain  substance  is  30  and  that 
of  another  is  60,  it  is  evident  that  30  grams  of  the  first  contain  as 
many  molecules  as  60  grams  of  the  latter.  Solutions  of  non- 
electrolytes  which  contain  in  the  same  amount  of  solvent  the 
weights  of  the  dissolved  substances  proportional  to  their  molecular 
weights,  have  the  same  boiling-point;  it  follows,  therefore,  that 
solutions  which  contain  the  same  number  of  molecules  in  the 
same  volume  boil  at  the  same  point.  With  these  facts  in  mind 
the  behavior  of  solutions  of  electrolytes  can  be  understood. 
According  to  the  electrolytic  dissociation  theory  molecules  of 
these  substances  break  down  into  ions;  a  solution  of  sodium 
chloride,  for  example,  if  completely  ionized,  contains  twice  as 


226  INORGANIC  CHEMISTRY  FOR  COLLEGES 

many  particles  in  solution  as  it  would  contain  if  it  were  not 
ionized : 

NaCl  ->  Na+  +  Cl~ 

If  we  assume  that  the  rise  in  the  boiling-point  of  a  solution  is  deter- 
mined not  by  the  number  of  molecules  in  solution,  but  by  the 
number  of  particles,  molecules  or  ions,  the  behavior  of  acids,  bases, 
and  salts  finds  a  ready  explanation.  If  dilute  solutions  of  sugar 
and  salt  are  made  by  dissolving  in  equal  amounts  of  water  a  gram- 
molecule  of  each,  it  will  be  found  that  the  rise  in  the  boiling-point 
of  the  salt  solution  is  just  twice  that  of  the  sugar  solution;  each 
molecule  of  salt  breaks  down  into  two  ions,  whereas  the  molecule  of 
sugar  does  not  change.  As  a  consequence,  one  solution  contains 
twice  as  many  particles  as  the  other  and  the  rise  in  its  boiling-point 
is  twice  as  great. 

When  1  gram-molecule  of  salt  is  dissolved  in  1000  c.c.  of  water 
the  rise  in  the  boiling-point  is  0.87°;  if  no  dissociation  had  taken 
place  the  rise  would  have  been  0.52°,  and  if  it  had  been  complete  it 
would  have  been  1.04°.  Evidently  but  a  part  of  the  molecules 
undergo  dissociation  at  this  concentration.  The  fraction  which 
breaks  down  into  ions  can  be  calculated  as  follows:  Consider  100 
molecules  and  let  a  equal  the  number  which  are  converted  into  ions ; 
a  is,  then,  the  percentage  dissociated.  The  number  of  un-ionized 
molecules  is  expressed  by  100  —  a;  and  since  in  the  case  of  sodium 
chloride  two  ions  are  formed  from  one  molecule,  2a  represents 
the  number  of  ions.  The  total  number  of  particles  formed  from  100 
molecules  is  then  100  —  a  +  2a,  which  simplified  is  100  +  a. 
Now,  since  the  rise  in  the  boiling-points  of  solutions  is  proportional 
to  the  number  of  particles  present,  we  have: 

0.52  ? 

~  0.87'  °] 

The  salt  is,  then,  dissociated  to  the  extent  of  67  per  cent. 

It  has  been  found  that  the  calculation  of  the  degree  of  dissocia- 
tion by  means  of  the  determination  of  the  rise  in  the  boiling- 
point  of  solutions  gives  results  which  agree  with  those  arrived  at 
by  other  methods  only  when  dilute  solutions  are  used.  When 
weights  greater  than  one-tenth  of  a  gram-molecule  are  dissolved  in  a 


ACIDS,    BASES,    SALTS.    SOLUTIONS  227 

liter  of  water  deviations  are  observed,  which  grow  greater  as  the 
concentration  is  increased. 

254.  The  Freezing-point  of  Solutions. — When  substances  are 
dissolved  in  water  and  other  liquids,  the  freezing-points  of  the 
solutions  are  lower  than  those  of  the  pure  solvents.     Laws  similar 
to  those  stated  above  in  regard  to  the  elevation  of  the  boiling-point 
have  been  found  to  express  the  influence  of  dissolved  substances 
on  the  freezing-points  of  solutions.     The  effect  on  the  freezing- 
point  is  greater  than  on  the  boiling-point.     A  molecular  weight 
of  a  non-electrolyte  dissolved  in  1000  grams  of  water  raises  the 
boiling-point  0.52°,  but  lowers  the  freezing-point  1.86°.    The  extent 
of  dissociation  can  be  calculated  from  the  freezing-point  in  the 
same  way  as  that  illustrated  above  in  the  case  of  the  boiling-point. 

When  dilute  aqueous  solutions  freeze,  pure  water  usually  sepa- 
rates out  in  the  solid  form.  In  order  that  the  ice  may  be  in  equi- 
librium with  the  solution,  the  temperature  must  be  that  at  which 
the  melting  of  the  ice  and  its  formation  from  the  solution  take 
place  at  the  same  rate.  This  temperature  is  below  0°. 

255.  A  saturated  solution  of  salt  freezes  at  —21°,  which  is  6° 
below  zero  Fahrenheit.     This  means  that  at  this  temperature  pure 
ice  and  a  saturated  solution  of  salt  are  in  equilibrium.     It  will  be 
recalled  that  the  freezing-point  is  denned  as  the  temperature  at 
which  the  solid  and  liquid  phases  are  in  equilibrium  (187),  and  that 
at  equilibrium  an  addition  of  heat  does  not  cause  a  rise  in  tempera- 
ture;  some  of  the  solid  changes  to  liquid.     If  solid  salt  is  put  on 
ice  at  0°,  some  of  the  former  dissolves  and  a  saturated  solution  of 
salt  is  formed;    the  amount  which  dissolves  may  be  vanishingly 
small,  but  the  solution  which  results  can  be  in  equilibrium  with  ice 
at  —  21  °  only.     In  order  that  the  temperature  may  fall,  some  of  the 
heat  present  in  the  ice  which  has  the  temperature  0°  must  dis- 
appear;  the  ice  melts  and  in  so  doing  absorbs  heat  and  the  tem- 
perature falls.     As  the  ice  melts  more  salt  dissolves  in  the  water, 
and  the  melting  continues  until  the  temperature  is  —21°.     These 
facts  are  utilized  in  making  freezing-mixtures  for  freezing  ice-cream 
and  in  keeping  switches  on  car  tracks  free  from  ice.     A  mixture  of 
100  parts  of  calcium  chloride,  CaCk,  GH^O  and  70  parts  of  snow 
will  lower  the  temperature  to  —54°. 

256.  Determination  of  Extent  of  lonization  from  the  Conduc- 
tivity of  Solutions. — The  extent  to  which  an  acid,  base,  or  salt 


228  INORGANIC  CHEMISTRY  FOR  COLLEGES 

ionizes  can  also  be  determined  by  a  study  of  the  conductivity  of 
solutions.  We  have  learned  that  the  best  explanation  which  has 
been  offered  of  the  way  in  which  a  solution  conducts  electricity 
is  based  upon  the  assumption  that  ions  are  present  in  the  solutions 
(243) .  If  this  is  true  the  larger  the  number  of  ions  present  the  more 
readily  will  the  solution  conduct.  The  extent  to  which  an  elec- 
trolyte breaks  down  into  ions  is  determined  by  the  amount  of 
water  present.  The  experiment  with  acetic  acid  (242)  showed 
that  the  acid  alone  did  not  conduct  the  current,  but  its  solution  in 
water  did  conduct,  and  that  the  conductivity  increased  more  and 
more  as  water  was  added  and  the  solution  was  diluted.  When  a 
similar  experiment  is  made  with  hydrochloric  acid  or  salt,  it  is  found 
that  a  point  is  finally  reached  when  the  addition  of  more  water  has 
no  effect  on  the  conductivity;  at  this  point  the  molecules  are 
assumed  to  be  completely  dissociated  into  ions.  In  order  to  find 
out  the  extent  to  which  a  salt  is  dissociated  into  ions  at  any  con- 
centration, we  first  determine  the  conductivity  of  its  solution,  when 
completely  ionized,  in  an  apparatus  arranged  so  that  all  the  ions 
take  part  in  carrying  the  current,  and  then  the  conductivity  of 
the  solution  having  the  special  concentration.  Suppose  in  the 
first  case  the  conductivity  is  represented  by  the  number  10  and  in 
the  next  case  by  6  then  0.6  of  the  salt  was  ionized  in  the  second 
case. 

257.  Methods  of  Expressing  Concentration  of  Solutions. — It 
has  been  stated  that  the  extent  of  the  ionization  of  electrolytes  is 
determined  by  the  relation  between  the  amounts  of  the  solvent  and 
the  dissolved  substance.  This  relation — the  concentration — can 
be  expressed  in  a  number  of  ways;  one  is  to  state  the  number  of 
grams  dissolved  in  a  liter  of  the  solution;  another  method,  which  is 
particularly  valuable  when  solutions  of  various  substances  are  to 
be  compared,  is  to  express  concentration  in  gram-molecular-weights 
per  liter.  A  solution  which  contains  1  gram-molecular-weight  of 
the  dissolved  substance  in  1  liter  of  the  solution  is  called  a  molar 
solution;  if  one-tenth  of  this  amount  is  present  the  solution  is  0.1 
molar.  When  acids  are  compared  another  method  is  often  used. 
The  characteristic  properties  of  acids  are  due  to  the  presence  of  the 
replaceable  hydrogen  atoms  they  contain.  A  molecule  of  hydro- 
chloric acid,  HC1,  is  not  equivalent  to  one  of  sulphuric  acid,  H2SO4, 
because  the  latter  contains  two  replaceable  hydrogen  atoms;  1 


ACIDS,    BASES,    SALTS.    SOLUTIONS  229 

gram-molecular-weight  of  sulphuric  acid  will  neutralize  twice  as 
much  sodium  hydroxide  as  1  gram-molecular-weight  of  hydro- 
chloric acid.  For  this  reason  another  method  is  used  in  expressing 
the  concentration  of  acids. 

A  normal  solution  of  an  acid  is  defined  as  one  which  contains 
in  1  liter  of  solution  the  amount  of  the  acid  that  furnishes 
1  gram  of  replaceable  hydrogen.  A  normal  solution  of  hydro- 
chloric acid  and  all  monobasic  acids  contains  1  gram-molecule  in  a 
liter.  A  normal  solution  of  sulphuric  acid  and  all  dibasic  acids 
contains  one-half  of  a  gram-molecule  in  a  liter.  A  normal  solution 
of  phosphoric  acid,  HsPCU,  contains  one-third  of  a  gram-molecular- 
weight  in  a  liter,  etc.  It  is  evident  that  a  definite  volume  of  a 
normal  solution  of  any  acid  is  equivalent  to  the  same  volume  of  a 
normal  solution  of  any  other  acid.  If  25  c.c.  of  a  normal  hydro- 
chloric acid  solution  is  required  to  react  with  a  certain  weight  of 
zinc  or  to  neutralize  a  definite  amount  of  sodium  hydroxide,  then 
25  c.c.  of  a  normal  solution  of  any  other  acid  will  react  with  these 
amounts. 

The  concentrations  of  solutions  of  bases  are  expressed  in  a 
similar  way.  A  normal  solution  of  a  base  which  has  one  hydroxyl 
group,  such  as  NaOH,  contains  1  gram-molecular-weight  in  a  liter. 
A  normal  solution  of  a  di-acid  base,  like  Ca(OH)2,  contains  one- 
half  of  a  gram-molecule  in  a  liter.  It  is  evident  that  1  c.c.  of  a 
normal  solution  of  any  base  will  neutralize  1  c.c.  of  a  normal  solu- 
tion of  any  acid. 

A  normal  solution  of  an  oxidizing  agent  contains  enough  of  the 
compound  dissolved  in  a  liter  to  oxidize  1  gram  of  hydrogen,  and 
a  normal  solution  of  a  reducing  agent  contains  in  a  liter  that 
amount  of  the  substance  which  equals  in  reducing  power  1  gram 
of  hydrogen.  The  concentration  of  any  solution  can  be  expressed 
as  a  fraction  or  multiple  of  a  normal  solution;  we  can  have,  for 
example,  a  twice  normal,  a  one-tenth  normal,  a  0.95  normal  and 
a  ^j  normal  solution  of  hydrochloric  acid,  etc.;  these  concentra- 
tions are  expressed  thus:  2N  HC1,  0.1N  HC1,  0.95N  HC1,  and 
N/64  HC1. 

Other  terms  are  used  in  expressing  the  concentrations  of  solu- 
tions when  reference  is  not  made  to  a  definite  amount  of  solvent 
or  solute;  thus  a  dilute  solution  contains  a  relatively  small  amount 
of  the  solute  as  compared  with  that  of  the  solvent;  a  concentrated 


230  INORGANIC  CHEMISTRY  FOR  COLLEGES 

solution  contains  a  relatively  large  amount  of  the  dissolved  sul> 
stance. 

A  saturated  solution  contains  as  much  of  the  dissolved  sub- 
stance as  can  be  held  in  solution  when  some  of  the  solid  substance 
is  present.  The  restriction  stated  in  the  last  clause  is  necessary, 
because  it  is  possible  for  a  liquid  to  hold  in  solution  more  of  a  solid 
than  can  be  presented  in  a  saturated  solution.  This  will  be  clear 
from  the  following  considerations.  The  solubility  of  most  salts 
increases  with  rise  in  temperature.  If  a  saturated  solution  of  such 
a  salt  is  made  at  a  certain  temperature,  and  is  then  separated 
from  any  solid  present,  it  is  often  possible  to  lower  the  tempera- 
ture and  have  all  the  salt  remain  in  solution.  Such  a  solution  is 
said  to  be  super-saturated,  for  if  a  bit  of  the  solid  is  added,  crystalli- 
zation takes  place  at  once  and  the  concentration  of  the  solution 
decreases  until  it  reaches  that  of  a  saturated  solution — when  the 
solution  and  the  solid  are  in  equilibrium. 

258.  The  Extent  of  the  lonization  of  Acids,  Bases,  and  Salts.— 
It  has  been  stated  that  the  characteristic  properties  of  solutions  of 
acids  are  due  to  the  hydrogen  ions  which  are  present.  If  this  is 
true  the  behavior  of  all  acids  should  be  in  general  alike,  but  there 
should  be  a  difference  in  activity  and  the  rate  at  which  they  react, 
if  the  acids  differ  markedly  in  the  extent  to  which  they  ionize. 
We  find  this  to  be  the  case.  If  we  place  pieces  of  zinc  of  equal 
size  and  shape  in  normal  solutions  of  hydrochloric,  sulphuric,  and 
acetic  acids,  we  find  that  the  rate  at  which  hydrogen  is  evolved  in 
the  three  cases  is  different.  A  determination  of  the  extent  to 
which  these  acids  ionize,  by  the  methods  outlined  above,  shows 
that  they  differ  markedly  in  this  respect. 

The  extent  to  which  an  electrolyte  ionizes  has  a  marked  effect 
on  its  chemical  behavior.  Acids  which  are  largely  ionized  are  often 
spoken  of  as  strong  acids  because  their  characteristic  acidic  proper- 
ties are  marked;  acids  that  ionize  but  slightly  are  weak  acids. 
For  example,  hydrochloric  acid  in  a  one-tenth  normal  solution  at 
18°  is  ionized  to  the  extent  of  92  per  cent.  Under  the  same  con- 
ditions of  concentration  and  temperature  sulphuric  acid  is  61  per 
cent  ionized,  and  nitric  acid  92  per  cent.  These  are  all  strong  acids. 
Carbonic  acid  in  one-twenty-fifth  normal  solution  is  ionized  to  the 
extent  of  0.2  per  cent  only;  it  is  a  very  weak  acid.  The  hydroxides 
of  sodium  and  potassium  are  strong  bases;  they  are  91  per  cent 


ACIDS,    BASES,    SALTS.    SOLUTIONS  231 

ionized  in  0.1N  solutions.  Calcium  and  barium  hydroxides  are 
also  strong  bases;  in  N/64  solutions  they  are  about  91  per  cent 
ionized.  Ammonium  hydroxide  is  a  relatively  weak  base;  it  is 
ionized  to  the  extent  of  1.3  per  cent  in  0.1N  solutions.  Most  salts 
are  largely  ionized.  The  percentages  of  ionization  of  a  few  salts 
in  0.1N  solutions  at  18°  are  as  follows:  Sodium  chloride  84,  sodium 
nitrate  83.  sodium  sulphate  70,  sodium  acetate  79.  The  values 
for  the  potassium  salts  are  about  the  same  as  those  of  the  sodium 
salts.  Copper  sulphate  and  zinc  sulphate  are  ionized  to  a  less 
extent — about  40  per  cent — and  are  typical  of  salts  which  yield 
bivalent  ions.  See  also  599,  page  511. 

EXERCISES 

1.  If  you  were  given  a  sample  of  an  element  what  experiments  would 
you  perform  to  determine  whether  it  was  an  acid-forming  or  a  base-forming 
element? 

2.  A  sample  of  hydrochloric  acid  was  neutralized  with  a  solution  of  sodium 
hydroxide  which  contained  30  grams  of  the  base  in  1  liter  of  solution.     In 
three  experiments  16.42,  16.50,  and  16.35  c.c.  of  the  base  were  used  to  neu- 
tralize 10  c.c.  of  the  acid.     Calculate  the  weight  of  hydrogen  chloride  in 
1000  c.c.  of  the  solution  of  the  acid. 

3.  Write  the  formulas  of  the  ions  formed  when  compounds  of  the  follow- 
ing composition  are  dissolved  in  water:    (a)  BaCl2,    (6)  KNO3,    (c)  NiSO4, 
(d}  Na,HP04,    (e)  CoSO4,    (/)  A1C13,    (0)  K2SO4,    (/*)  NaHCO3,    (t)  K2CO3, 
0")  Ca(N03)2,   (/c)  K2Cr04,    (0  NaHSO3. 

4.  The  compounds  having  the  composition  represented  by  the  following 
formulas  are  insoluble:    (a)  BaSO4,    (6)  CaCO3,    (c)  Ca3(PO4)2,    (d)  AgBr, 
(e)  CuS,    (/)  A1(OH)3,  (g)  Cu(OH)2.     Write  equations,  using  ionic  formulas 
for  reactions,  by  which  each  compound  can  be  prepared  by  double  decompo- 
sition. 

5.  Write  ionic  equations  for  the  following:    (a)  the  formation  of  ferrous 
chloride,  FeCl2,  by  the  action  of  hydrochloric  acid  on  iron;    (6)  the  elec- 
trolysis of  a  solution  of  salt;    (c)  the  action  of  iron  on  a  solution  of  copper 
sulphate,    (d)  the  neutralization  of  calcium  hydroxide  by  nitric  acid. 

6.  An  experiment  showed  that  when  3  grams  of  a  certain  non-electrolyte 
were  dissolved  in  100  c.c.  of  water  the  boiling-point  of  the  solution  was  100.26°. 
Calculate    (a)  the  molecular  weight  of  the  compound  and    (6)  the  freezing- 
point  of  the  solution. 

7.  A  solution  of  1.5  grams  of  a  non-electrolyte  in  50  c.c.  of  water  was  found 
to  freeze  at  —  0.6°.     Calculate    (a)  the  molecular  weight  of  the  compound 
and   (6)  the  boiling-point  of  the  solution. 

8.  When  0.1  gram-molecular-weight  of  an  acid  that  formed  two  ions  was 
dissolved  in  1000  c.c.  of  water,  the  freezing-point  of  the  solution  was  found 
to  be  —  0.357°.     Calculate  the  percentage  dissociation  of  the  acid. 


232  INORGANIC  CHEMISTRY  FOR  COLLEGES 

9.  What  weights  of  the  following  must  be  dissolved  in  enough  water  to 
make  1000  c.c.  of  solution  in  preparing  molar  solutions  (a)  HC1,  (6)  H2SO4, 
(c)  KOH,   (d)  Ba(OH)2,  (e)  BaCl2,   (/)  CuSO4,5H2O?     (g)  What  is  the  nor- 
mality of  the  solutions  of  the  acids  and  bases? 

10.  Express  the  normality  of  solutions  that  contain  (a)  42  grams  HC1 
per  liter,    (6)  23  grams  H2SO4  per  liter,   (c)  10  grams  Ba(OH)2  per  liter. 

11.  It  was  found  that  1.6  c.c.  of  a  solution  of  hydrochloric  acid  were  required 
to  neutralize  1  c.c.  of  a  normal  solution  of  sodium  hydroxide.      What  is  the 
normality  of  the  acid  solution? 


CHAPTER  XVIII 
CHEMICAL  EQUILIBRIUM 

259.  Chemical  reactions  take  place  between  molecules,  and 
these  are  physical  entities;  as  a  consequence,  for  a  complete  inter- 
pretation of  such  reactions  we  must  take  into  account  not  only 
the  substances  involved  and  their  chemical  properties,  but  the 
physical  conditions  under  which  the  molecules  are  brought  together, 
such  as  concentration  and  temperature.  If  we  limit  the  dis- 
cussion for  the  present  to  gases,  it  is  clear  that  since  chemical 
reaction  can  take  place  only  when  the  molecules  approach  one 
another  and  come  into  what  might  be  called  chemical  contact, 
the  concentration  of  the  molecules  is  an  important  factor  in 
the  rate  at  which  they  interact.  If  in  a  given  volume  we  have  a 
given  number  of  two  kinds  of  molecules  which  react  with  each 
other  at  a  certain  rate,  it  is  evident  that  if  we  introduce  into  this 
volume  a  larger  number  of  these  molecules  the  number  of  con- 
tacts between  them  in  a  given  time  will  increase,  and  the  rate  of  the 
interaction  will  be  greater— more  molecules  will  be  converted  into 
the  product  of  the  reaction.  Again,  if  we  raise  the  temperature  of 
the  molecules  they  will  move  more  rapidly,  their  contacts  will  be 
more  frequent,  as  a  consequence,  and  the  rate  of  the  reaction  will 
increase.  We  see,  therefore,  that  concentration  and  temperature 
—physical  conditions — are  important  factors  in  determining  the 
rate  at  which  molecules  interact.  Up  to  the  present  the  chemical 
aspects  of  the  reactions  studied  have  been  emphasized, and  attention 
has  been  centered  on  the  products  formed  as  the  result  of  these 
reactions;  but  we  are  interested  also  in  how  these  transformations 
come  about — the  mechanics  of  molecular  action  in  chemical  change. 
The  detailed  study  of  chemistry  from  this  point  of  view  has  led  to 
generalizations  of  great  value,  which  have  not  only  broadened 
markedly  our  knowledge  of  chemical  phenomena,  but,  as  we  shall 

233 


234  INORGANIC  CHEMISTRY  FOR  COLLEGES 

see,  have  been  utilized  in  solving  many  problems  of  the  greatest 
technical  importance. 

260.  Reversible  Reactions.  —  Many  chemical  reactions  are 
reversible,  that  is,  the  substances  formed  as  the  result  of  the  reac- 
tion are  able,  under  the  proper  circumstances,  to  interact  and  pro- 
duce the  substances  from  which  they  were  formed.  This  can  be 
expressed  by  "the  following  equation  : 


where  A  and  B  represent  elements  or  compounds  which  through 
the  expenditure  of  chemical  energy  pass  into  the  substances  rep- 
resented by  C  and  D.  If  the  latter  can  react  to  form  A  and  B  the 
reaction  is  said  to  be  reversible. 

261.  Chemical  Equilibrium.  —  It  is  evident  that  when  a  reaction 
of  the  above  type  takes  place  all  four  substances  will  be  present. 
A  and  B  react  at  a  definite  rate  which  is  determined  by  their 
mutual  chemical  affinities;  likewise  C  and  D,  but  the  rate  in  this 
case  is  different.     When  the  reaction  starts  the  amounts  of  A  and  B 
will  decrease  and  those  of  C  and  D  increase,  but  a  time  will  come, 
as  the  reaction  proceeds,  when  the  quantity  of  any  one  substance 
formed  just  equals  the  quantity  of  it  which  undergoes  trans- 
formation.    When  this  condition  has  been  reached  the  reaction  is 
said  to  be  in  equilibrium.     It  is  important  to  emphazise  the  fact 
that  chemists  are  of  the  opinion  that  the  molecules  are  in  a  state 
of    constant    change  —  the    equilibrium    is    dynamic,    not    static. 
This  type  of  equilibrium  has  been  emphasized  already  in  con- 
nection with  vaporization  (179)  .     When  such  a  state  of  equilibrium 
exists  as  the  result  of  the  continued  change  of  molecules  into  other 
molecules  of  a  different  kind  it  is  defined  as  chemical  equilibrium. 
An  example  of  a  reaction  which  reaches  an  equilibrium  has  already 
been  given  (125): 

C12  +  H2O  <=±  HOC1  +  HC1 

In  this  case  equilibrium  results  in  a  solution  of  1  volume  of  chlorine 
in  1  of  water  when  66  per  cent  of  the  chlorine  is  present  as  molecules 
of  the  element  and  34  per  cent  as  hypochlorous  and  hydrochloric 
acids. 

262.  Effect  of  Concentration  on    Equilibrium.  —  The  concen- 
tration of  the  substances  in  a  chemical  reaction  at  equilibrium  can 


CHEMICAL  EQUILIBRIUM  235 

be  varied  by  adding  more  of  one  of  the  substances  or  by  taking 
away  some  of  one.  In  a  typical  reaction  of  this  type, 

A  +  B<=±C  +  D 

if  more  of  A  is  added  and  the  volume  kept  constant  so  that  the  con- 
centration of  A  is  increased,  there  will  be  a  larger  number  of  con- 
tacts between  A  and  B  in  a  given  time;  as  a  consequence,  the  rate 
at  which  C  and  D  are  produced  will  be  greater,  and  their  amounts 
present  when  equilibrium  is  reached  will  be  increased.  It  follows, 
therefore,  as  a  consequence  of  increasing  the  concentration  of  one 
or  more  of  the  substances  taking  part  in  a  reversible  reaction,  that 
a  change  takes  place  as  a  result  of  which  more  of  the  products  pro- 
duced from  these  substances  are  present  when  equilibrium  is 
attained.  This  fact  expressed  more  definitely  and  in  a  mathe- 
matical form  is  known  as  the  law  of  molecular  concentration,  which 
is  sometimes  called  the  law  of  mass  action  (see  264).  An  exam- 
ple will  serve  to  emphasize  this  important  conclusion.  At  high 
temperatures  the  reaction  expressed  by  the  following  equation  is 
reversible : 

2SO2  +  O2  <=±  2SO3 

Sulphur  dioxide  reacts  with  oxygen  to  form  sulphur  trioxide,  but 
at  the  same  temperature  the  trioxide  decomposes  into  sulphur 
dioxide  and  oxygen.  It  has  been  found  by  experiment  that  when 
the  dioxide  and  oxygen  are  mixed  in  the  proportion  represented  by 
the  equation  and  heated  to  720°  about  40  per  cent  of  sulphur  di- 
oxide is  converted  into  trioxide.  If,  however,  twice  as  much 
oxygen  is  used  and  the  volume  of  the  gases  kept  the  same  as  before 
so  that  the  concentration  of  the  oxygen  is  doubled,  about  60  per 
cent  sulphur  dioxide  is  changed. 

The  important  consequences  of  the  change  in  the  equilibrium 
with  change  in  concentration  are  evident.  If  we  are  preparing  a 
substance  by  means  of  a  reaction  which  is  reversible — and  this  is 
often  necessary — we  can  increase  the  amount  of  product  formed,  by 
increasing  the  concentration  of  one  or  more  of  the  reacting  sub- 
stances. The  excess  of  the  material  which  is  present  when  equi- 
librium is  reached  can  be  recovered;  its  presence  during  the  reac- 
tion was  important  since  it  effected  the  result  attained,  but  it  was 
not  used  up. 


236  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  concentrations  of  the  substances  taking  part  in  a  rever- 
sible reaction  can  be  changed  by  taking  away  one  of  the  products 
of  the  reaction.  If,  for  example,  in  the  reaction  represented  by  the 
equation 


C  is  removed  as  fast  as  it  is  formed,  it  is  evident  that  there  is  no 
opportunity  for  C  and  D  to  produce  A  and  B.  As  a  consequence, 
the  reaction  runs  to  completion  and  A  and  B  are  completely  trans- 
formed into  C  and  D.  An  important  example  of  this  kind  of  a 
reaction  is  what  we  have  called  double  decompositions  (148). 
When,  for  example,  sodium  chloride  is  treated  with  sulphuric  acid 
the  reaction  represented  by  the  following  equation  takes  place  : 

NaCl  +  H2SO4  =  NaHSO4  +  HC1 

In  this  case  the  hydrogen  chloride  is  a  gas  and  escapes  and  cannot, 
therefore,  interact  with  the  sodium  hydrogen  sulphate  to  form 
sodium  chloride.  The  reaction  is  a  reversible  one,  however;  if  it 
is  carried  out  in  such  a  way  that  the  hydrogen  chloride  cannot 
escape,  all  of  the  salt  is  not  changed  to  sodium  hydrogen  sulphate. 
In  fact,  by  selecting  the  proper  physical  conditions,  we  can  con- 
vert sodium  hydrogen  sulphate  and  hydrogen  chloride  into  sodium 
chloride  and  sulphuric  acid  and  thus  completely  reverse  the  reac- 
tion. This  can  be  done  by  adding  the  sulphate  to  a  concentrated 
solution  of  hydrochloric  acid.  In  this  case  the  hydrogen  chloride 
is  kept  in  the  field  of  reaction  by  being  dissolved  in  water.  Sodium 
chloride  is  insoluble  in  concentrated  hydrochloric  acid  and  pre- 
cipitates. The  reaction  in  this  case  is  expressed  by  the  above 
equation  written  backward: 

NaHSO4  +  HC1  =  NaCl  +  H2SO4 

A  double  decomposition  takes  place,  since  the  sodium  chloride  is 
automatically  removed  as  the  result  of  the  fact  that  it  is  insoluble 
in  the  liquid  present.  The  high  concentration  of  the  hydrogen 
chloride  in  the  saturated  solution  is  a  factor  in  forcing  the  equi- 
librium to  completion  as  indicated  by  the  equation. 

From  the  standpoint  of  chemical  equilibrium  we  can  now  under- 
stand the  reason  for  the  statement  which  has  been  given  as  to  the 
conditions  under  which  a  reversible  reaction  becomes  one  of  double 


CHEMICAL  EQUILIBRIUM  237 

decomposition.  This  occurs,  it  will  be  recalled,  when,  as  the  prod- 
uct of  the  reaction,  there  is  formed  a  gas,  a  substance  which  decom- 
poses into  a  gas,  or  an  insoluble  substance.  In  these  cases  the 
product  which  possesses  any  one  of  the  above  properties  is  auto- 
matically removed  from  the  sphere  of  action.  Double  decom- 
positions which  result  from  the  formation  of  undissociated 
substances  are  explained  as  follows : 

In  solution,  most  inorganic  reactions  take  place  between  ions, 
and  these  play  the  same  part  as  the  molecules  do  when  reactions 
occur  between  the  latter.  For  example,  in  a  reaction  of  neu- 
tralization, such  as  that  between  sodium  hydroxide  and  hydro- 
chloric acid,  the  theories  accepted  at  present  lead  to  the  view  that 
the  reaction  takes  place  between  ions  which  are  represented  in  the 
following  equation: 

Na+  +  OH~  +  H+  +  Cr  =  Na+  +  Cl"  +  H2O 

Since  the  sodium  ion  and  the  chlorine  ion  do  not  take  part  in  the 
reaction,  the  latter  can  be  simplified  to  the  following : 

H+  +  OH~  ±=>  H2O 

The  equilibrium  in  this  reaction  is  such  that  practically  all  of  the 
ions  change  into  undissociated  water;  as  the  result,  the  action  of  an 
acid  and  a  base  is  one  of  double  decomposition. 

263.  Equilibrium  in  Homogeneous  and  in  Heterogeneous  Sys- 
tems.— In  the  discussion  up  to  this  point  it  has  been  assumed  that 
all  of  the  substances  involved  in  the  reactions  are  present  in  one 
phase  (188) — that  is,  they  are  uniformly  distributed.  In  this  case 
the  system  is  said  to  be  homogeneous.  If  two  or  more  of  the  sub- 
stances are  present  in  different  phases,  for  example,  if  one  is  a  gas 
and  one  is  a  solid,  the  system  is  said  to  be  heterogeneous.  If  we 
are  considering  an  equilibrium  in  the  gaseous  or  liquid  phase  in 
such  a  system,  the  concentration  of  the  solid  in  these  phases  is 
constant  as  long  as  any  of  the  solid  is  present.  For  example,  in 
the  reversible  reaction  represented  by  the  equation 

4H2O  +  3Fe  *±  Fe3O4  +  4H2 

the  relation  between  the  amounts  of  water  and  hydrogen  present  at 
equilibrium  is  not  changed  by  altering  the  amounts  of  iron  or  iron 
oxide,  provided  some  of  each  is  present. 


238  INORGANIC  CHEMISTRY  FOR  COLLEGES 

264.  The  Law  of  Molecular  Concentration.—  The  effect  of 
concentration  on  chemical  equilibrium  can  be  put  into  a  simple 
mathematical  form  which  makes  it  possible  to  calculate  how  much 
of  each  of  the  reacting  substances  is  present  at  equilibrium  at  any 
concentration,  provided  the  amounts  present  at  one  concentration 
are  known.  It  is  necessary  to  find  these  amounts  for  one  set  of 
concentrations  by  experiment.  In  the  reversible  reaction  repre- 
sented as  follows  : 


the  rate  at  which  A  and  B  are  transformed  into  C  and  D  is  pro- 
portional to  the  molecular  concentration  of  A  and  B  and  to  the 
mutual  chemical  affinities  between  the  latter.  Likewise,  the  rate 
at  which  C  and  D  are  transformed  into  A  and  B  is  proportional 
to  the  chemical  affinity  which  exists  between  C  and  D  and  to  their 
molecular  concentrations.  Since  the  chemical  affinities  are  constant 
at  a  fixed  temperature,  the  rates  of  the  two  reactions  under  these 
conditions  are  proportional  to  the  concentrations  only.  It  will  be 
recalled  (note,  section  88)  that  if  a  quantity  varies  as  two  other 
quantities  vary,  the  first  will  be  proportional  to  the  product  of 
the  latter;  and  that  if  one  quantity  is  proportional  to  another, 
the  first  is  equal  to  a  constant  multiplied  by  the  second.  Making 
use  of  these  mathematical  modes  of  expression,  we  can  put  the 
statement  in  regard  to  the  effect  of  concentration  on  equilibrium 
into  the  form  of  a  formula.  The  rate,  R'  of  the  reaction 
A  +  B  =C  +  Dis  proportional  to  the  molecular  concentration 
of  A,(CA)  and  of  B3(CB);  if  K'  represents  a  constant,  then 

R'  =  K'  X  CA  X  CB 

The  similar  equation  for  the  opposed  reaction  C-r-E)  =A  +  Bis 
R"  =  K"  X  Cc  X  CD 

At  equilibrium  these  rates  are  equal,  and,  therefore,  under  these 
circumstances  R'  =  R"  and 

K'  X  CA  X  CB  =  K"  X  Cc  X  CD, 
or 

C>C    X    OD  -K-  -rr- 

CAXCB  =K77  =  K 


CHEMICAL  EQUILIBRIUM  239 

It  is  customary  in  writing  equations  to  express  equilibria,  to  put  in 
the  numerator  of  the  fraction  the  molecular  concentrations  of  the 
products  represented  on  the  right  side  of  the  chemical  equation  for 
the  reaction  involved.  The  last  equation  given  above  states  that 
in  a  reversible  reaction  the  product  of  the  molecular  concentra- 
tions of  the  substances  formed,  divided  by  the  product  of  the  molec- 
ular concentrations  of  the  interacting  substances,  is  a  constant. 

This  constant  is  known  as  the  equilibrium  constant  of  the 
reaction.  The  general  statement  given  above  is  one  mode  of 
expression  of  the  so-called  law  of  molecular  concentration.  It  was 
proposed  in  1864  by  Guldberg  and  Waage,  two  Norwegian  chem- 
ists, and  was  based  on  their  own  experiments  as  well  as  those  of 
Wilhelmy,  who  had  studied  the  influence  of  concentration  on  the 
rate  of  reactions,  and  those  of  Gladstone,  who  had  investigated 
the  influence  of  concentration  on  chemical  equilibrium.  The  law 
was  originally  called  the  mass  law  but  the  name  was  not  well  chosen, 
since  it  does  not  apply  to  mass  but  to  molecular  concentration. 
For  example,  in  using  the  law  in  connection  with  a  particular  re- 
action we  do  not  express  the  concentration  in  grams,  by  which 
mass  is  measured,  but  in  moles;  the  law  was  deduced  as  the  result 
of  the  consideration  of  the  effect  on  the  equilibrium  of  an  in- 
crease in  the  number  of  molecules  taking  part  in  any  reaction. 

A  knowledge  of  the  value  of  the  equilibrium  constant  is  of  the 
greatest  importance  when  a  reversible  reaction  is  being  used  for 
the  preparation  of  any  of  the  substances  involved,  for  it  makes  it 
possible  to  calculate  to  what  extent  changes  in  concentration  will 
effect  the  quantity  of  the  desired  substance  obtained. 

If  a  reversible  reaction  is  of  the  type  represented  by  the  equa- 
tion 

A  +  2B  =  C  +  D, 

the  mathematical  expression  of  the  law  takes  a  slightly  different 
form.  In  this  case  two  molecules  of  B  react  with  one  of  A,  or 
A  +  B  +  B=C  +  D;  accordingly,  R'  =  K'  X  CA  X  CB  X  CB 
or  R'  =  K'  — »  CA  X  CB2.  It  is  seen  from  this  that  the  most  gen- 
eral expression  of  the  law  is  as  follows : 

CpU  X  Cpv  .  .  .    _   -p,- 

CA"  X  CB*  •  •  •  = 


240  INORGANIC  CHEMISTRY  FOR  COLLEGES 

where  u,  v,  w,  and  x  represent,  respectively,  the  number  of  mole- 
cules of  the  several  substances  which  enter  into  the  reaction. 

265.  Effect  of  Temperature  on  the  Rate  of  Chemical  Reac- 
tions.— Since  the  equilibrium  attained  in  a  reversible  reaction  is 
dependent  on  the  relative  rates  at  which  the  two  opposing  reac- 
tions take  place,  it  is  necessary  to  consider  the  effect  of  change  in 
temperature  on  these  rates.     It  has  been  found  by  experiment 
that  within  ordinary  ranges  of  temperature  the  rate  at  which  reac- 
tions take  place  is  approximately  doubled  for  a  rise  of  10  degrees. 
From  the  standpoint  of  the  kinetic  theory  this  fact  appears  reason- 
able, for  with  rise  in  temperature  it  is  postulated  that  the  molecules 
move  faster  and  come  into  contact  more  frequently;   as  a  conse- 
quence, more  of  the  products  of  the  reaction  are  formed  in  a  given 
time — the  rate  of  the  reaction  increases. 

The  fact  that  reactions  proceed  more  rapidly  as  the  tempera- 
ture is  increased  is  utilized  by  the  chemist.  A  reaction  proceeding 
at  a  certain  rate  at  room  temperature  (20°)  will  go  approximately 
256  times  as  fast  at  100°.  The  change  in  temperature  in  this 
case  is  80  degrees,  and  since  the  rate  is  doubled  for  each  10  degrees, 
the  rate  is  28  =  256  times  as  fast  at  the  higher  temperature. 
Many  inorganic  reactions,  such  as  double  decompositions,  take 
place  practically  instantaneously  and  heating  is  unnecessary.  When 
reactions  take  place  slowly,  however,  such  as  the  decomposition  of 
potassium  chlorate  into  potassium  chloride  and  oxygen,  they  are 
usually  carried  out  at  more  or  less  elevated  temperatures. 

266.  Effect  of  Temperature  on  Chemical  Equilibrium. — When 
a  system  in  chemical  equilibrium  is  heated,  the  rates  of  the  two 
opposing  reactions  are  both  increased.     This  change  results  from 
the  increased  motion  of  the  molecules  and  the  change  in  the  chem- 
ical affinities  between  the  interacting  substances.     If  the  tempera- 
ture of  any  definite  system  in  equilibrium  is  raised,  the  effect  of  the 
increase  is  the  same  on  the  motion  of  all  the  molecules,  and  the 
change  in  the  rates  of  the  opposing  reactions  due  to  this  cause  alone 
will  be  the  same.     But  increase  in  temperature  changes  the  inten- 
sity of  the  chemical  affinities  of  all  molecules  differently;    conse- 
quently the  change  in  the  rates  at  which  the  two  reactions  proceed 
will  be  different  when  the  temperature  of  the  system  is  raised,  and 
the  concentrations  at  equilibrium  will,  as  a  result,  be  different. 

The  effect  of  temperature  on  chemical  equilibrium  has  been 


CHEMICAL  EQUILIBRIUM  241 

extensively  investigated  on  account  of  its  importance.  It  was 
found  to  be  associated,  as  one  might  expect,  with  the  amounts  of 
heat  evolved  or  absorbed  when  chemical  reactions  occur.  If 
we  take,  for  example,  the  reaction  represented  by  the  following 
equation, 

A  +  B<=±C-f-D  +  z  cal. 

when  A  and  B  react  to  form  C  and  D,  heat  to  the  amount  of  x 
calories  is  set  free;  on  the  other  hand  when  C  and  D  react  to  form 
A  and  B  the  same  amount  of  heat  is  absorbed — transformed  into 
chemical  energy.  The  change  of  heat  energy  into  chemical  energy 
takes  place  more  readily  at  high  than  at  low  temperatures;  as  a 
consequence,  when  the  temperature  of  a  system  represented  by  the 
equation  given  above  is  raised,  the  rate  of  the  reaction  by  which  C 
and  D  change  to  A  and  B  is  increased  and  the  equilibrium  shifts 
so  that  more  of  the  latter  are  present.  The  equilibrium  is  shifted 
from  right  to  left  in  the  above  equation  by  increasing  the  concen- 
tration of  C  or  of  D  or  the  temperature  at  which  the  reaction  is  car- 
ried out.  It  is  evident  that  if  we  are  using  such  a  reaction  for  the 
preparation  of  C  or  D,  it  should  be  carried  out  at  the  lowest  tem- 
perature possible,  at  which  the  rate  is  satisfactory.  Since  the  rate 
of  the  reaction  may  be  very  small  at  the  temperature  which  yields 
a  satisfactory  concentration  of  the  desired  substance,  such  reac- 
tions are  often  brought  about  in  the  presence  of  a  catalyst.  It 
will  be  recalled  that  catalysts  markedly  increase  the  rate  of  reac- 
tions. 

In  the  case  of  an  endothermic  reaction  such  as, 

E  +  F  —  G  +  H-z  cal. 

rise  in  temperature  leads  to  the  shifting  of  the  equilibrium  from 
left  to  right.  In  this  case  when  E  and  F  change  into  G  and  H  heat 
is  absorbed,  and  as  heat  can  be  more  readily  changed  into  chemical 
energy  at  high  temperatures,  rise  in  temperature  favors  the  forma- 
tion of  G  and  H.  The  direction  in  which  the  equilibrium  shifts 
can  also  be  seen  by  changing  the  equation  to  read, 

E  +  F  +  x  cal.  <=»  G  +  H 

and  using  the  method  employed  in  the  previous  example;  it  is  seen 
that  increase  in  temperature  shifts  the  equilibrium  so  that  more 
G  and  H  are  formed. 


242  INORGANIC  CHEMISTRY  FOR  COLLEGES 

It  follows  from  the  above  that  if  a  substance  is  to  be  prepared 
through  the  use  of  an  endothermic  reaction,  the  preparation  should 
be  carried  out  at  as  high  a  temperature  as  possible.  Under  these 
conditions  the  rate  of  the  reaction  is  usually  high  and  no  catalyst 
is  required. 

In  general,  when  a  large  amount  of  heat  is  evolved  or  absorbed 
in  a  reaction,  the  effect  of  change  in  temperature  on  the  equilibrium 
is  great;  on  the  other  hand, when  the  heat  change  is  small  the  effect 
of  change  in  temperature  is  small.  A  knowledge  of  this  principle 
is  of  great  value  in  coming  to  a  conclusion  as  to  whether  a  par- 
ticular reaction  could  be  used  for  the  preparation  on  the  industrial 
scale  of  one  of  the  products  of  the  reaction.  For  example,  if  the 
reaction  is  an  endothermic  one  and  the  equilibrium  at  a  definite 
temperature  is  such  that  but  a  small  proportion  of  the  desired 
product  is  formed,  whether  or  not  the  reaction  can  be  used  will 
depend,  among  other  considerations,  on  the  heat  exchange.  If 
this  is  great,  the  proportion  of  the  desired  product  at  high  tempera- 
tures will  be  much  increased— enough  perhaps  to  make  it  advisable 
to  use  the  reaction.  If  the  heat  change  is  small,  increase  in  tem- 
perature will  have  little  effect  on  the  equilibrium. 

267.  The  Law  of  van't  Hoff. — The  effect  of  temperature  on 
equilibrium  was  studied  by  van't  Hoff,  a  Dutch  chemist,  who 
generalized  the  facts  in  regard  to  it  into  a  form  which  is  known  as 
van't  Hoff's  law  of  mobile  equilibrium.  This  law  states  that 
when  the  temperature  of  a  system  in  equilibrium  is  raised,  the 
equilibrium-point  is  displaced  in  the  direction  which  absorbs  heat. 
This  law  is  of  the  greatest  importance  and  along  with  the  law  of 
molecular  concentration  is  the  guiding  principle  followed  in  the 
study  and  utilization  of  all  reversible  reactions. 

The  law  of  van't  Hoff  applies  to  physical  equilibria,  such,  for 
example,  as  that  existing  between  a  salt  and  its  saturated  solution. 
If  we  represent  this  equilibrium  as  follows: 

Solid  salt  +  water  ^±  solution  of  salt  —  x  cal. 

the  fact  is  indicated  that  when  the  salt  dissolves,  heat  is  absorbed. 
Applying  van't  Hoff's  law  we  would  expect  that  with  rise  in  tem- 
perature the  solubility  would  increase — and  this  is  the  fact.  Anhy- 
drous sodium  sulphate  is  an  example  of  another  type  of  salt; 
when  it  dissolves  heat  is  given  off.  In  this  case  the  solubility 
decreases  with  rise  in  temperature. 


CHEMICAL  EQUILIBRIUM  243 

When  sodium  chloride,  common  salt,  dissolves  in  water  the 
heat  change  is  small  (  —  1,180  cal.),  and  the  effect  of  rise  in  tem- 
perature on  its  solubility  is  small.  On  the  other  hand,  when  potas- 
sium nitrate  dissolves,  a  much  larger  amount  of  heat  is  absorbed 
( —8,520  cal.  per  mol) ;  this  results  in  the  fact  that  the  temperature 
of  the  solution  drops  when  a  strong  solution  of  the  salt  is  made. 
In  this  case  the  solubility  of  the  salt  is  markedly  affected  by  rise 
in  temperature. 

268.  The  Law  of  Le  Chatelier. — The  law  of  van't  Hoff  applies 
to  certain  cases  coming  under  the  broader  generalization  first  put 
forward  by  Le  Chatelier.     This  law,  which  applies  to  both  physical 
and  chemical  equilibria,  sums  up  the  behavior  of  substances  when 
they  are  subjected  to  a  change  in  conditions.     It  is  as  follows:  If 
some  stress  is  brought  to  bear  on  a  system  in   equilibrium,  the 
equilibrium  is  displaced  in  such  a  direction  that  the  normal  effect 
of  the  stress  results.1     By  stress  is  meant  the  effect  of  any  thing 
or  any  kind  of  energy  which  has  an  effect  on  the  equilibrium;   it 
may  be,  for  example,  change  in  concentration,  pressure,  or  tem- 
perature. 

269.  The  Effect  of  Pressure  on  Equilibrium. — When  a  chemical 
reaction  takes  place  which  is  accompanied  by  a  change  in  volume, 
the  pressure  under  which  the  change  occurs  has  a  marked  effect 
on  the  equilibrium.     If  we  consider  the  reaction  by  which  ammonia 
is  formed  from  hydrogen  and  nitrogen, 

N2  +  3H2  ^  2NH3 

we  see  that  1  volume  of  nitrogen  unites  with  3  volumes  of  hydrogen 
to  form  2  volumes  of  ammonia  gas — there  is  a  contraction  from  4  to 
2  volumes.  By  making  use  of  Le  Chatelier's  law  we  can  discover 
in  which  direction  the  equilibrium  will  shift  if  the  pressure  on  the 
system  is  increased.  If  we  replace  the  words  "  some  stress  "  in 
the  above  statement  of  the  law  by  the  word  pressure,  the  law  will 
read, — if  pressure  is  brought  to  bear  on  a  system  in  equilibrium 
the  equilibrium  is  displaced  in  such  a  direction  that  the  normal 
effect  of  the  pressure  results.  Increase  in  pressure  brings  about  a 
decrease  in  volume;  as  a  consequence,  in  the  above  .reaction  under 
increased  pressure  more  ammonia  will  be  present  in  the  system  at 

1  This  law  is  usually  expressed  as  follows :  If  some  stress  is  brought  to 
bear  on  a  system  in  equilibrium,  the  equilibrium  is  displaced  in  the  direction 
which  tends  to  undo  the  effect  of  the  stress. 


244  INORGANIC  CHEMISTRY  FOR  COLLEGES 

equilibrium  because  when  the  change  of  nitrogen  and  hydrogen 
to  ammonia  takes  place  the  volume  decreases.  This  particular 
reaction  will  be  considered  in  detail  later  (339),  and  we  shall  see 
that  a  knowledge  of  the  facts  summarized  in  Le  Chatelier's  law 
made  it  possible  to  develop  a  process  for  the  manufacture  of  am- 
monia. 

EXERCISES 

1.  Salts  having  the  following  formulas  are  practically  insoluble  in  water: 
(a)  Agl,    (6)  BaCO3,    (c)  CaCO3,    (d)  CuS.     Write  equations  for  reactions 
by  which  each  can  be  formed,  and  state  how  the  principle  of  chemical  equi- 
librium applies  to  the  reactions. 

2.  Zinc  sulphide,   ZnS,   is  slightly  soluble  in   water,  but  its  solubility 
increases  in  the  presence  of  hydrochloric  acid,  the  amount  passing  into  solu- 
tion being  determined  by  the  concentration  of  the  acid.     State  reasons  for 
what  you  would  expect  to  happen  if  a  solution  of    (a)  Na2S  and  one  of    (b) 
H2S  were  added  to  a  solution  of  ZnCl2.     Copper  sulphide,  CuS,  is  insoluble 
in  hydrochloric  acid,     (c)  What  would  happen  if  solutions  a  and  b  above 
were  added  to  separate  portions  of  a  solution  of  copper  chloride,  CuCl2? 

3.  Nitric  acid  boils  at  86°.     Under  what  conditions  could  it  be  prepared 
by  double  decomposition  from  sodium  nitrate,  NaNO3,  and  sulphuric  acid? 

4.  Discuss  the  changes  which  occur  in  the  following  systems  when  heat 
is  applied:    (a)  H2O  (liquid)  +±  H2O  (gas);    (b)  H2O  (solid)  <=±  H2O  (liquid); 
(c)  Discuss  the  changes  which  occur  in  the  two  cases  when  the  temperature 
is  constant  but  the  pressure  is  increased. 

5.  Discuss  the  changes  produced  in  the  reversible  reaction  PC13  -f  C13 
«=^  PCU  +  33,000  cal.  when    (a)  the  temperature  is  changed  and    (b)  the 
pressure  is  changed,     (c)  If  PC16  is  converted  into  vapor  in  a  space  filled  with 
C12  how  would  its  dissociation  compare  in  extent  with  that  produced  when 
it  is  vaporized  in  the  same  space  filled  with  air? 

6.  Show  how  Henry's  law  is  consistent  with  LeChatelier's  law. 

7.  Would  you  expect  pressure  to  have  much  effect  on  the  solubility  of 
a  salt  in  water?     Give  a  reason  for  your  answer.     Why  does  the  temperature 
have  a  marked  effect  on  the  solubility  of  salts? 

8.  Lead  iodide,  which  is  slightly  soluble  in  water,  can  be  prepared  by  the 
following     reaction:      Pb(NO3)2+  2KI  =  PbI2  +  2KNO3.     (a)  Write     the 
reaction,  using  ionic  symbols,  and  state  what  would  happen  if  a  small  amount 
of  a  solution  of  lead  nitrate  were  added  to  a  saturated  aqueous  solution  of 
lead  iodide?     (b)  How  could  the  same  result  be  produced  without  the  addi- 
tion of  a  lead  salt? 

9.  Silver  chloride,  AgCl,  which  is  slightly  soluble  in  water,  is  frequently 
precipitated  and  weighed  in  making  a  quantitative  determination  of  chlorides. 
When  it  is  washed  to  remove  impurities  a  small  amount  of  silver  nitrate  is 
added  to  the  wash-water.     State  a  reason  for  the  use  of  the  latter. 


CHAPTER  XIX 
SULPHUR  AND  HYDROGEN  SULPHIDE 

270.  Sulphur  occurs  in  the  free  condition  in  volcanic  regions 
and  has  been  known  since  the  earliest  times.     It  is  mentioned  in 
the  Bible  and  in  Homer,  and  was  considered  to  be  one  of  the  ele- 
mentary principles  of  which  the  earth  is  composed.     It  represented 
the  spirit  of  fire  and  was  called  brimstone  (fire-stone).     In  the 
fifteenth  century  when  the  action  of  chemical  substances  on  the 
body  was  first  studied  in  an  endeavor  to  find  some  way  to  combat 
disease,  sulphur  was  used  as  a  medicine.     It  was  believed  that 
the  body  was  made  up  of  the  elementary  principles  mercury,  sul- 
phur, and  salt,  and  a  lack  in  one  of  these  produced  illness.     In  the 
case  of  certain  ailments,  sulphur  was  prescribed  and  in  the  case 
of  others,  compounds  of  mercury  or  various  kinds  of  salts.     As  the 
result  of  the  foundation  of  a  system  of  medicine  on  this  hypothesis, 
the  effect  on  the  body  of  a  great  many  chemical  substances  was 
discovered,  and  a  basis  was  laid  for  that  branch  of  modern  medicine 
which  is  called  pharmacology.     The  utilization  of  sulphur  as  a 
medicine  has  extended  to  modern  times,  although  at  present  its 
application  is  limited  to  use  in  ointments  and  salves  to  combat 
certain  diseases  of  the  skin.     Sulphur  is  a  necessary  constituent  of 
the  body,  and  it  must  be  present  in  our  food,  but  it  has  been  shown 
that  the  free  element  is  not  assimilated;  it  must  be  present  in  com- 
plicated organic  compounds  to  be  of  value,  and  it  is  from  these, 
such  as  the  protein  of  eggs,  meat,  etc.,  that  we  obtain  our  supply  of 
the  element.     Any  excess  of  sulphur  over  that  required  is  excreted 
as  sulphates  in  the  urine. 

271.  Occurrence  of  Sulphur. — Free  sulphur  is  found  abundantly 
in  Sicily,  mixed  with  limestone,  CaCOa,  gypsum,  CaSCU,  2H2O 
and  other  minerals.     It  also  occurs  in  Japan,  Italy,  Spain,  Cali- 
fornia, Egypt,  China,  and  India.     Sulphur  is  found  in  rock  deposits 
where  it  has  been  probably  formed  from  gypsum  through  the  action 

245 


246  INORGANIC  CHEMISTRY  FOR  COLLEGES 

of  bacteria.      Large  quantities  of  sulphur  are  found  in  the  sedi- 
mentary deposits  of  Texas  and  Louisiana. 

The  element  also  occurs  in  compounds;  hydrogen  sulphide, 
H2S,  which  is  a  gas  with  a  disagreeable  and  characteristic  odor, 
issues  from  volcanoes,  and  occurs  dissolved  in  the  water  of  sulphur 
springs;  sulphur  dioxide,  SOo,  is  also  a  product  of  volcanic  activity. 
Many  sulphides  are  important  minerals,  some  of  which  are  used  as  a 
source  of  sulphur  and  some  of  the  other  elements  which  they  con- 
tain; pyrite,  FeS2,  furnishes  sulphur  for  the  manufacture  of  sul- 
phuricacid ;  zincblende,  ZnS,  cinnabar, HgS, galena,  PbS,  and chalco- 
cite,  Cu2S,  are  valuable  ores.  Many  sulphates  occur  in  nature,  of 
which  gypsum,  CaSC>4,  2H2O,  is  perhaps  the  most  important. 

Sulphur  occurs  in  all  living  things;  it  is  a  constituent  of  certain 
proteins  and  is  found  in  hair,  nails,  meat,  eggs,  etc.  When  these 
compounds  undergo  putrefaction  the  sulphur  which  they  contain  is 
converted  into  hydrogen  sulphide  and  other  substances  which  pos- 
sess a  disagreeable  odor.  The  element  occurs  in  many  oils  found 
in  plants,  which  give  the  latter  their  most  characteristic  properties ; 
such  oils  have  been  extracted  from  onions,  mustard,  garlic,  cab- 
bage, etc. 

272.  Extraction  of  Sulphur. — The  chief  sources  of  sulphur  are 
Sicily  and  Louisiana.  In  Sicily  the  rock  with  which  the  sulphur  is 
mixed  is  piled  in  a  crude  kiln  and  covered  with  some  of  the  ore  left 
from  a  previous  burning.  The  kiln  is  fired  by  igniting  some  of  the 
sulphur,  and  allowing  it  to  burn  for  some  time;  the  draught  holes 
are  then  closed.  The  heat  generated  melts  the  sulphur,  which 
collects  in  the  bottom  of  the  kiln  and  is  drawn  off  from  time  to  time 
and  run  into  wooden  molds.  This  method  is  crude  and  wasteful; 
about  one-quarter  of  the  sulphur  burns  to  sulphur  dioxide,  which 
passes  into  the  air,  and  as  a  result  damages  vegetation  in  the 
neighborhood  of  the  kiln. 

The  sulphur  in  Louisiana,  which  is  found  in  deposits  over  100 
feet  thick  at  a  depth  of  about  450  feet,  is  obtained  by  an  ingenious 
method  devised  by  Frasch.  Four  concentric  pipes,  1,  4,  6,  and  10 
inches  in  diameter  were  sunk  into  the  deposit.  Between  the 
6-  and  10-inch  pipes  superheated  water  at  about  170°  is  pumped 
into  the  sulphur,  which  melts  at  this  temperature.  Hot  air  under 
pressure  is  forced  down  the  1-inch  pipe,  and,  as  a  result,  a  mix- 
ture of  air,  molten  sulphur,  and  water  rises  in  the  space  between  the 


SULPHUR  AND  HYDROGEN  SULPHIDE  247 

4-inch  and  6-inch  pipes.  The  material  is  run  into  wooden  tanks 
where  the  sulphur  solidifies.  The  product  is  quite  pure  and  is 
used  without  further  refining. 

Crude  sulphur  is  refined  by  distilling  it  from  iron  retorts  and 
conducting  the  vapor  produced  into  chambers  made  of  brick.  If 
the  temperature  of  the  condensing  chamber  is  below  110°  the  sul- 
phur collects  in  the  form  of  a  fine  powder,  called  "  flowers  of  sul- 
phur"; at  a  higher  temperature  the  vapor  condenses  to  a  liquid, 
which  is  run  off  into  cylindrical  molds,  where  it  solidifies.  In  this 
form  the  sulphur  is  often  called  roll  sulphur,  or  roll  brimstone. 

273.  Physical  Properties  of  Sulphur. — It  will  be  recalled  that 
carbon  can  exist  in  three  distinct  allotropic  forms — diamond, 
graphite,  and  amorphous  carbon.  Sulphur  shows  the  property  of 
existing  in  two  crystalline  allotropic  modifications,  but  in  this  case 
but  one  form  is  stable  at  room  temperature  and  the  other  slowly 
changes  to  it  on  standing.  The  stable  form  of  sulphur,  which 
occurs  native,  crystallizes  in  the  rhombic  system  and  is  called, 
therefore,  rhombic  sulphur.  It  is  pale  yellow  in  color,  is  brittle, 
odorless,  tasteless,  has  the  specific  gravity  2.06  and  melts  at  114.5°; 
it  is  almost  insoluble  in  water,  but  dissolves  in  carbon  disulphide 
(41  parts  in  100  at  18°),  forming  a  solution  from  which  it  can  be 
obtained  in  well-defined  crystals  by  evaporation. 

When  rhombic  sulphur  is  melted  and  allowed  to  cool,  the 
crystals  formed  are  long,  transparent  needles  which  belong  to  the 
monoclinic  system.  These  can  be  obtained  by  melting  some  sul- 
phur in  a  crucible,  and  before  it  has  completely  solidified,  pouring 
off  the  liquid  through  a  hole  made  in  the  solid  crust  formed  on  the 
surface  of  the  liquid.  Monoclinic  sulphur  is  almost  colorless,  melts 
at  119.25°,  has  the  specific  gravity  1.96,  and  dissolves  in  carbon 
disulphide.  It  slowly  changes  at  room  temperature  to  rhombic 
sulphur,  and  the  transparent  needles  become  opaque  as  the  result 
of  the  formation  of  minute  rhombic  crystals.  The  two  solid  forms 
of  sulphur  are  in  equilibrium  at  96°,  which  is  the  transition  point 
for  the  two  forms.  Below  96°  rhombic  sulphur  is  the  stable  form. 

When  either  form  of  sulphur  is  melted,  a  clear,  pale-yellow, 
limpid  liquid  is  first  obtained;  as  the  temperature  is  raised  the 
color  changes  to  dark  brown  and  the  liquid  becomes  at  about  160° 
so  viscous  that  it  will  scarcely  flow  out  of  the  vessel  containing  it, 
when  it  is  inverted.  As  the  temperature  is  increased  above  260° 


248  INORGANIC  CHEMISTRY  FOR  COLLEGES 

the  mixture  becomes  less  viscous,  and  at  444.7°  the  liquid  boils. 
The  yellow  liquid  is  represented  by  the  symbol  SX  and  the  brown 
liquid  by  Sju.  If  S/z  is  allowed  to  cool  slowly  it  passes  into  SX 
and  the  crystals  obtained  on  solidification  dissolve  in  carbon  disul- 
phide.  If,  however,  the  brown  molten  sulphur  is  poured  into 
water  and  chilled  suddenly  in  this  way,  the  sulphur  assumes  the 
form  of  a  plastic  mass  which  is  somewhat  elastic.  After  standing 
some  days  plastic  sulphur  becomes  hard  and  opaque;  it  consists 
of  a  mixture  of  rhombic  sulphur  and  S/x  in  the  form  of  an  amorphous 
solid;  the  two  forms  can  be  separated  by  carbon  disulphide  in 
which  S/i  is  insoluble.  SM  changes  very  slowly  to  rhombic  sulphur 
at  room  temperature,  a  number  of  years  being  necessary  to  effect 
complete  transformation. 

274.  Chemical  Properties  of  Sulphur. — Sulphur  is  an  active 
element  and  unites  with  both  metallic  and  non-metallic  elements. 
In  this  respect  it  resembles  oxygen  and  many  of  the  formulas  of 
the  sulphides  are  analogous  to  those  of  the  oxides.  This  is  due  to 
the  fact  that  when  sulphur  unites  with  metals  it  has  the  valence 
2,  which  is  the  valence  of  oxygen.  The  formulas  of  some  com- 
pounds produced  as  the  result  of  the  direct  union  of  sulphur  with 
other  elements  are  as  follows:  H2S,  ZnS,  CuS,  FeS,  Ag2S,  P2S5, 
and  082;  the  formulas  suggest  those  of  the  oxides  of  these  ele- 
ments. In  compounds  containing  the  more  negative  elements 
sulphur  exhibits  the  valence  4  or  6;  thus,  when  it  burns  sulphur 
dioxide,  SO2,  is  formed  along  with  a  small  amount  of  sulphur  tri- 
oxide,  SOa;  with  chlorine,  sulphur  forms  compounds  of  the  com- 
position S2C12,  SC12,  and  SCU.  Sulphur  monochloride,  S2C12,  is  a 
reddish-yellow  liquid,  boiling  at  138°,  which  dissolves  sulphur; 
the  solution  is  used  in  vulcanizing  rubber.  Sulphur  combines 
slowly  at  room  temperature  with  all  metals  except  the  least  active 
ones  like  platinum  and  gold.  Even  silver  which  is  not  affected  by 
oxygen  reacts  readily  with  sulphur.  If  a  bit  of  rubber  is  left  in 
contact  with  silver  the  latter  turns  black  as  the  result  of  a  reac- 
tion between  the  metal  and  the  sulphur  present  in  the  vulcanized 
rubber.  Silver  spoons  are  tarnished  when  left  in  contact  with 
eggs,  because  the  latter  contain  sulphur  compounds  which  are 
decomposed  as  the  result  of  the  affinity  of  silver  for  sulphur. 

Many  metals  burn  when  introduced  into  sulphur  vapor.  The 
union  of  iron  and  sulphur  has  already  been  described  (16). 


SULPHUR  AND  HYDROGEN  SULPHIDE  249 

Carbon  combines  with  sulphur,  but  as  the  reaction  is  an  endo- 
thermic  one,  energy  must  be  supplied  to  bring  about  the  formation 
of  carbon  disulphide  (215). 

275.  Uses  of  Sulphur. — In  the  crude  condition  sulphur  is  used 
for  making  sulphur  dioxide,  which,  in  turn,  is  employed  in  the 
manufacture  of  sulphuric  acid,  in  bleaching,  and  for  many  other 
purposes.     Large  quantities  of  flowers  of  sulphur  are  consumed  in 
destroying  a  fungus  which  causes  a  disease  in  grapevines;  the  value 
of  the  element  for  this  purpose  is  probably  due  to  the  fact  that  in 
the  presence  of  the  oxygen  and  water  in  the  air  it  slowly  changes  to 
sulphuric  acid,  which  prevents  the  growth  of  the  fungus.     Refined 
sulphur  is  used  in  the  manufacture  of  matches,  gunpowder,  fire- 
works, and  for  vulcanizing  rubber.     Many  of  the  important  newer 
dyes  are  made  by  heating  with  free  sulphur  certain  compounds 
prepared  from  the  products  found  in  coal-tar;  some  of  these  dyes 
contain  traces  of  free  sulphur  or  compounds  which  give  up  the 
element  readily,  and,  as  a  consequence,  blacken  silver  which  is 
wrapped  in  material  dyed  with  them.     The  industrial  importance 
of  sulphur  can  be  seen  from  the  fact  that  over  800,000  tons  of  it 
are  used  annually. 

HYDROGEN  SULPHIDE 

276.  Hydrogen  sulphide,  H2$,  is  found  in  the  gases  which  issue 
from  volcanoes  and  in  the  water  of  sulphur  springs.     It  is  formed 
in  the  putrefaction  of  certain  proteins,  which  are  the  most  impor- 
tant nitrogenous  constituents  of  animal  and  vegetable  substances. 
The  disagreeable  odor  produced  as  the  result  of  the  decomposition 
is  in  part  due  to  this  gas.     Rotten  eggs  are  said  to  smell  of  hydrogen 
sulphide,  but  an  egg  can  have  undergone  decomposition  and  possess 
a  most  disagreeable  odor  before  a  test  will  show  the  presence  of 
the  gas. 

277.  Preparation  of  Hydrogen  Sulphide. — Hydrogen  sulphide 
can  be  made  by  passing  hydrogen  through  boiling  sulphur,  but  the 
method  lacks  practical  significance: 

H2  H-  S  =  H2S 

It  is  prepared  in  the  laboratory  by  the  action  of  hydrochloric  acid 
on  ferrous  sulphide : 

FeS  +  2HC1  =  FeCl2  +  H2S 


250  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  reaction  is  similar  to  that  between  ferrous  oxide  and  hydro- 
chloric acid: 

FeO  +  2HC1  =  FeCl2  +  H2O 

It  has  been  pointed  out  that  sulphur  resembles  oxygen  in  many  of 
its  reactions;  the  behavior  of  sulphides  in  certain  cases  is  similar 
to  that  of  oxides.  There  are  differences,  however;  every  metal 
forms  an  oxide  which  will  react  with  hydrochloric  acid,  but  all 
sulphides  do  not  dissolve  in  the  acid.  The  sulphides  of  the 
metals  from  potassium  to  iron  inclusive,  in  the  electro-motive 
series  of  the  elements  (252),  dissolve  in  dilute  hydrochloric  acid. 
Any  of  these  could  be  used  to  make  hydrogen  sulphide,  but  iron 
sulphide  is  commonly  employed  because  it  is  readily  prepared  from 
iron  and  sulphur.  Other  acids  can  replace  hydrochloric  acid,  but 
nitric  acid  can  not,  because  it  oxidizes  hydrogen  sulphide. 

When  many  organic  substances  which  contain  hydrogen  and 
sulphur  are  heated  to  a  high  temperature  or  are  burned  in  an 
insufficient  supply  of  air,  hydrogen  sulphide  is  formed.  A  decom- 
position of  this  kind  leads  to  the  formation  of  hydrogen  sulphide 
when  coal  is  heated  to  prepare  coal-gas,  or  when  it  is  burned  with 
insufficient  draft  to  produce  complete  combustion  (223). 

278.  Physical  Properties  of  Hydrogen  Sulphide. — Hydrogen 
sulphide  is  a  colorless  gas,  somewhat  heavier  than  air;    1  volume 
of  water  dissolves  4.37  volumes  of  hydrogen  sulphide  at  0°;  3.60  at 
10°,  and  3.23  at  15°;   the  gas  is  insoluble  in  boiling  water.     On 
account  of  the  solubility  of  hydrogen  sulphide  in  water  it  is  usually 
collected  in  the  laboratory  by  upward  displacement  of  air.     The 
critical  temperature  of  the  gas  is  100°;  below  this  it  can  be  lique- 
fied by  pressure.     Hydrogen  sulphide  boils  at  —62°  and  melts  at 
—83°.     The  gas  is  a  powerful  poison;    it  produces  nausea  and 
coma  and  1  part  in  200  of  air  causes  death. 

279.  Heat  of  Formation  of  Hydrogen  Sulphide. — When  hydro- 
gen unites  with  sulphur  to  form  hydrogen  sulphide  a  comparatively 
small  amount  of  energy  is  set  free.     The  thermo-chemical  equa- 
tion when  solid  sulphur  is  used  is  as  follows: 

H2  +  S  =  H2S  +  4,800  cals. 

In  the  formation  of  water  much  more  energy  is  liberated;  when  2 
grams  of  hydrogen  burn  in  oxygen  58,100  calories  are  liberated. 


SULPHUR  AND  HYDROGEN  SULPHIDE  251 

We  should  expect,  therefore,  to  find  hydrogen  sulphide  a  much 
less  stable  compound  than  water,  and  being  unstable,  more  react- 
ive. When  a  large  amount  of  chemical  energy  is  transformed 
into  heat  in  the  formation  of  a  compound,  the  compound  is 
usually  stable  when  heated. 

280.  Chemical  Behavior  of  Hydrogen  Sulphide. — With  the 
above  facts  in  mind  we  are  not  surprised  to  find  that  hydrogen 
sulphide  begins  to  dissociate  into  hydrogen  and  sulphur  at  a  com- 
paratively low  temperature,  310°,  as  compared  with  that  at  which 
water  begins  to  dissociate,  about  1800°. 

Hydrogen  sulphide  burns  in  air  with  a  pale  blue  flame : 

2H2S  +  3O2  =  2H2O  +  2SO2 

In  this  case  a  large  amount  of  energy  can  be  changed  into  heat  as 
the  result  of  the  union  of  hydrogen  and  sulphur  with  oxygen; 
when  1  gram-molecular-weight  of  hydrogen  sulphide  is  burned 
122,500  calories  are  set  free.  In  general,  compounds  which  are 
combustible  are  composed  of  elements  that  burn. 

At  room  temperature  a  solution  of  hydrogen  sulphide  is  decom- 
posed slowly  by  the  oxygen  in  the  air.  Under  these  conditions  the 
hydrogen  is  oxidized  to  water  and  sulphur  separates  as  a  white 
powder: 

2H2S  +  02  =  2H20  +  2S 

A  similar  reaction  takes  place  when  hydrogen  sulphide  is  burned 
in  a  limited  supply  of  air.  All  the  metals  down  to  and  including 
silver  in  the  electromotive  series  of  the  metals  decompose  hydrogen 
sulphide  in  the  cold,  and  hydrogen  and  the  sulphide  of  the  metal  are 
formed.  The  tarnishing  of  metals  in  air  containing  hydrogen  sul- 
phide is  due  to  this  cause.  If  an  exceedingly  thin  layer  of  sulphide 
is  formed  on  silver  it  has  a  yellow  or  golden  tint,  but  as  the  amount 
of  sulphide  increases  the  color  changes  to  black. 

Hydrogen  sulphide  reacts  with  oxygen  not  only  when  it  is  in 
the  free  condition  but  when  it  is  combined  with  other  elements. 
For  example,  it  reacts  with  sulphur  dioxide  according  to  the  fol- 
lowing equation: 

2H2S  +  SO2  =  2H2O  +  3S 

We  see  again  in  this  reaction  the  strong  tendency  of  hydrogen  to 
unite  with  oxygen.  It  is  probable  that  the  sulphur  found  near 


252  INORGANIC  CHEMISTRY  FOR  COLLEGES 

volcanoes  is  produced  as  the  result  of  this  reaction.  When  hydrogen 
sulphide  reacts  in  this  way  it  is  said  to  be  a  reducing  agent,  for  it 
takes  oxygen  away  from  the  sulphur;  as  a  result  of  the  reaction 
the  sulphur  dioxide  is  reduced  and  the  hydrogen  sulphide  is  oxi- 
dized. 

A  solution  of  hydrogen  sulphide  in  water  shows  an  acid  reac- 
tion with  litmus,  but  it  is  a  very  weak  acid  as  it  is  dissociated  in 
one-tenth  normal  solution  to  the  extent  of  0.07  per  cent  only.  It 
ionizes  in  steps  as  other  dibasic  acids  do: 

H2S  ->  H+  +  HS~ 


The  ionization  according  to  the  second  expression  is  very  slight, 
being  less  than  that  of  water.  When  hydrogen  sulphide  is  passed 
into  a  solution  of  sodium  hydroxide,  sodium  hydrogen  sulphide, 
NaHS,  is  formed  : 

Na+  +  OH~  +  H+  +  HS~  =  Na+  +  HS~  +  H2O 

The  solution  of  the  salt  is  neutral  on  account  of  the  fact  that 
practically  no  hydrogen  ions  are  present. 

281.  If  a  solution  of  sodium  hydrogen  sulphide  is  mixed  with 
one  of  sodium  hydroxide  and  the  water  is  driven  off  by  heat,  the 
normal  salt  is  obtained  : 

NaOH  +  NaHS  <=>  Na2S  +  H2O 

The  reaction  is  a  reversible  one;  if  sodium  sulphide  is  dissolved  in 
water  the  reaction  indicated  by  reading  the  above  equation  from 
right  to  left  takes  place.  Sodium  hydroxide  and  sodium  hydrogen 
sulphide  are  formed,  and  the  solution  shows  a  strongly  alkaline 
reaction.  In  this  case  hydrolysis  is  said  to  have  taken  place. 
Many  substances  are  hydrolyzed  by  water  and  reactions  of  this 
kind  are  of  importance.  The  change  which  takes  place  when  water 
reacts  with  a  compound  and  converts  it  into  two  or  more  compounds 
is  called  hydrolysis.  The  action  of  an  acid  and  a  base  to  form  a 
salt  and  water  is  called  neutralization  (239);  the  reverse  of  this, 
the  action  of  water  with  a  salt  to  form  an  acid  and  a  base  is  an 
example  of  hydrolysis.  In  some  cases,  as  in  the  above,  an  acid 
salt  is  formed.  The  cause  of  hydrolysis  is  clear  from  a  considera- 


SULPHUR  AND  HYDROGEN  SULPHIDE  253 

tion  of  the  ions  involved  in  the  reaction.     The  equation  for  the 
action  of  water  on  sodium  sulphide  is  as  follows  : 

2Na+  +  S~     +  H+  +  OH~  =  2Na+  +  HS~  +  OH~ 


Water  breaks  down  to  a  very  small  extent  into  H+  and  OH~  ions, 
the  H4"  from  the  water  and  the  S~~  from  the  sulphide  unite  to 
form  HS~,  and,  as  a  result,  OH  ions  are  left  in  the  solution  which, 
accordingly,  shows  an  alkaline  reaction.  See  section  600,  page  511. 
282.  Sulphides.  —  All  the  metals  form  sulphides;  the  com- 
pounds differ  in  solubility  so  widely  that  one  of  the  most  impor- 
tant parts  of  qualitative  chemical  analysis  involves  the  preparation 
and  separation  of  sulphides.  The  behavior  of  solutions  of  the 
salts  of  the  various  metals  when  hydrogen  sulphide  is  passed 
through  them  can  be  illustrated  by  a  few  typical  examples.  The 
gas  is  passed  successively  into  solutions  of  the  following  salts, 
sodium  chloride,  calcium  chloride,  zinc  sulphate,  ferrous  chloride, 
copper  sulphate,  and  arsenic  chloride.  Nothing  takes  place  in  the 
first  two  cases;  in  the  others,  precipitates  are  produced  as  the 
result  of  the  formation  of  the  sulphides  of  the  elements.  The 
equations  for  the  reactions  are  as  follows: 

ZnSO4  +  H2S  =  ZnS  +  H2SO4 

FeCls  +  H2S  =  FeS  +  2HC1 

CuSO4  +  H2S  =  CuS  +  H2SO4 

2AsCl3  +  3H2S  =  As2S3  +  6HC1 

The  sulphide  of  zinc  is  white,  that  of  arsenic  yellow,  while  ferrous 
sulphide  and  copper  sulphide  are  black;  the  color  of  the  sulphides 
is,  thus,  helpful  in  their  identification.  If  dilute  hydrochloric  acid 
is  added  to  the  four  sulphides,  only  those  of  zinc  and  iron  dissolve; 
these  sulphides  are  not  precipitated  if  the  solution  is  acidified  before 
hydrogen  sulphide  is  introduced,  because  they  are  soluble  in  acids. 
The  reactions  indicated  by  the  first  and  second  equations  above 
are  reversible.  Hydrogen  sulphide  does  not  completely  precipitate 
these  sulphides,  because  as  a  result  of  the  reaction  an  acid  is 
produced.  If  it  is  desired  to  have  the  reaction  proceed  to  comple- 
tion, a  salt  of  hydrogen  sulphide  must  be  used,  and  for  this  pur- 
pose ammonium  sulphide  is  generally  employed: 

ZnSO4  +  (NH4)2S  =  ZnS  +  (NH4)2SO4 


254  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Under  these  circumstances  no  acid  is  formed  and  the  reaction  runs 
to  completion. 

The  behavior  of  the  salts  of  the  metals  illustrated  above  is 
typical  of  many  others;  some  are  not  precipitated  by  hydrogen 
sulphide  from  an  aqueous  solution  (potassium  to  aluminium 
inclusive  in  the  electromotive  series),  some  are  not  precipitated  in 
the  presence  of  acids  but  are  precipitated  by  ammonium  sulphide 
(manganese  to  iron  inclusive),  and  some  are  precipitated  by  hydro- 
gen sulphide  in  the  presence  of  dilute  acids  (cobalt  to  gold  inclu- 
sive). Of  the  sulphides  iasoluble  in  dilute  acids  those  of  tin,  arsenic, 
antimony,  platinum,  and  gold  are  soluble  in  yellow  ammonium 
sulphide,  (NH^S,.  By  means  of  the  reactions  outlined  above 
the  metals  can  be  separated  into  various  groups  (see  Appendix  VI) 
and  by  making  use  of  the  action  of  other  reagents  the  metals 
in  a  single  group  can  be  separated  and  identified.  These  tests 
will  be  given  later  when  the  metals  are  described  separately. 

283.  Test  for  Sulphides. — Hydrogen  sulphide  is  recognized  by 
its  characteristic  odor  or  by  the  fact  that  it  produces  a  coloration 
when  it  comes  into  contact  with  a  piece  of  paper  moistened  with  a 
solution  of  a  lead  salt;  lead  acetate  is  usually  used  in  making  the 
test.  The  equation  for  the  reaction  is — 

H2S  +  Pb(C2H3O2)2  =  PbS  +  2H(C2H3O2) 

The  color  of  the  lead  sulphide  precipitated  varies  from  yellow  to 
black  according  to  the  amount  of  gas  present. 

In  testing  for  a  sulphide  the  substance  is  warmed  with  dilute 
hydrochloric  acid,  and  if  a  gas  is  given  off  it  is  tested  by  exposing 
it  to  a  piece  of  moist  lead  acetate  paper.  Many  sulphides  do  not 
react  with  dilute  hydrochloric  acid.  In  testing  for  these  a  small 
piece  of  zinc  is  added  to  the  mixture  of  acid  and  the  substance 
under  examination.  If  a  sulphide  insoluble  in  acids  is  present  it  is 
reduced  by  the  nascent  hydrogen,  and  hydrogen  sulphide  is  formed. 
The  equations  for  the  reactions  in  the  case  of  copper  sulphide  are  as 
follows: 

Zn  +  2HC1  =  ZnCl2  +  2H 

CuS  +  2H  =  Cu  +  H2S 
The  gas  given  off  is  tested  as  before  with  lead  acetate  paper. 


SULPHUR  AND  HYDROGEN  SULPHIDE  255 

284.  Polysulphides. — Sulphur  dissolves  in  solutions  of  sodium 
sulphide  and  forms  a  mixture  of  compounds  to  which  the  formulas 
Na2§2,  Na2$3,  and  Na2Ss  have  been  assigned.  If  the  solution 
containing  these  salts  is  poured  into  a  dilute  acicj,  an  oil  is  obtained 
from  which  liquid  sulphides  of  hydrogen  have  been  isolated.  The 
one  having  the  formula  11282  is  unstable  and  resembles  somewhat 
hydrogen  peroxide,  H2O2,  in  properties.  Of  compounds  of  this 
type,  which  are  known  as  polysulphides,  the  one  prepared  by  dis- 
solving sulphur  in  ammonium  sulphide  is  used  in  qualitative  analy- 
sis. Its  formula  is  usually  written  (NH^S*,  the  x  indicating 
that  the  substance  is  a  mixture  of  polysulphides.  Sodium  poly- 
sulphide  is  used  in  the  manufacture  of  sulphur  dyes,  and  calcium 
poly  sulphide,  as  a  fungicide. 

EXERCISES 

1.  Calculate  the  percentage    by  weight  of  hydrogen  sulphide  contained 
in  a  solution  of  the  gas  in  water  saturated  at  0°. 

2.  What  weight  of  ferrous  sulphide  which  is  90  per  cent  pure  is  required 
to  make  enough  hydrogen  sulphide  to  fill  a  tank  holding  5  cu.  ft.  with  the  gas 
at  4  atmospheres  pressure  at  0°? 

3.  By  what  chemical  reactions  could  you  convert  (a)  zinc  sulphide  into 
zinc  oxide,    (6)  copper  oxide  into  copper  sulphide,    (c)  ferrous  sulphide  into 
ferrous  sulphate,    (d)  copper  sulphide  into  copper,    (e)  lead  into  lead  sul- 
phide? 

4.  Brass  contains  copper  and  zinc.     How  could  the  metals   (a)  be  brought 
into  solution,   and     (6)  separated  through  the  use  of  hydrogen  sulphide? 
(c)  How  could  the  free  metals  be  obtained  from  the  sulphides? 


CHAPTER  XX 
THE  OXIDES  AND  ACIDS  OF  SULPHUR 

285.  On  account  of  the  fact  that  sulphur  burns  and  occurs  free 
in  nature,  sulphur  dioxide  has  been  known  from  the  earliest  times. 
It  was  used  in  certain  religious  ceremonies  by  the  Greeks,  and 
its  disagreeable  properties,  no  doubt,  led  to  the  selection  of  fire  and 
brimstone  in  depicting  the  horrors  of  hell.     Priestley  first  isolated 
the  gas  in  pure  condition  as  the  result  of  an  accident.     He  was 
attempting  to  distill  sulphuric  acid  when  his  thermometer,  which 
contained   mercury,    broke,    and    sulphur   dioxide   was    formed. 
Priestley  in  his  writings  emphasizes  the  fact  that  many  of  his  most 
important  discoveries  were  made  as  the  result  of  accident,  but  it 
should  be  noted  that  he  had  the  ability  to  discern  the  important 
result  of  the  accident  and  interpret  it.     Lavoisier  explained  the 
behavior  of  sulphur  dioxide  with  water  in  1777,  and  indicated  the 
relation  between  the  acid  formed  and  sulphuric  acid  by  naming  the 
former  sulphurous  acid. 

286.  Occurrence  of  Sulphur  Dioxide. — This  compound  is  found 
in  the  gases  which  issue  from  volcanoes.     Great  quantities  of  it  are 
introduced  into  the  air  as  the  result  of  the  burning  of  coal  which 
contains  sulphur.     The  gas  which  causes  the  disagreeable  choking 
sensation  in  poorly  ventilated  railroad  stations  is  sulphur  dioxide. 
An  important  ore  of  copper  is  a  sulphide;   when  this  is  smelted 
the  sulphur  is  gotten  rid  of  by  heating  it  in  air,  and  large  quantities 
of  sulphur  dioxide  are  formed.     As  a  result,  in  the  neighborhood  of 
smelters  vegetation  is  destroyed  for  miles.     Laws  have  been  passed 
which  prohibit  the  introduction  into  the  air  of  the  products  of  the 
furnaces  which  contain  more  than  a  definite  amount  of  sulphur 
dioxide.     In  order  to  comply  with  these  requirements  it  was  neces- 
sary to  find  some  way  to  dispose  of  the  obnoxious  gas.     A  profit- 
able solution  of  the  problem  was  found  in  converting  it  into  sul- 
phuric acid,  which  is  a  valuable  product;    in  one  of  the  largest 

256 


THE  OXIDES  AND  ACIDS  OF  SULPHUR  257 

sulphuric  acid  plants  in  the  world  the  source  of  the  sulphur  is  the 
waste  gases  obtained  in  the  extraction  of  copper  from  a  sulphide 
ore. 

287.  Preparation  of  Sulphur  Dioxide.  —  For  industrial  purposes 
sulphur  dioxide  is  prepared  by  burning  sulphur  or  heating  a  sul- 
phide in  air.  Iron  pyrites,  FeS2,  is  generally  used  for  this  purpose: 

4FeS2  +  11O2  =  2Fe2O3  +  8SO2 

Sulphur  dioxide  can  be  conveniently  prepared  by  the  action 
of  an  acid  on  a  sulphite,  which  is  a  salt  of  sulphurous  acid,  H2S03. 
When  sodium  sulphite  and  hydrochloric  acid  are  used  the  reactions 
represented  by  the  following  equations  take  place  : 

Na2SO3  +  2HC1  =  2NaCl  +  H2SO3 
H2SO3=H2O 


The  reaction  resembles  that  between  sodium  carbonate,  Na2COs, 
and  the  acid;  in  each  case  an  acid  is  formed  which  breaks  down 
into  its  anhydride  and  water.  In  the  laboratory  a  saturated  solu- 
tion of  sodium  bisulphite  is  commonly  used,  as  it  can  be  bought 
at  a  low  price,  and  concentrated  sulphuric  acid  is  allowed  to  run 
into  it  slowly  from  a  dropping  funnel.  The  equation  for  the  reac- 
tion is  as  follows  : 

NaHSO3  +  H2SO4  =  NaHSO4  +  SO2  +  H2O 

Sulphur  dioxide  can  also  be  prepared  by  reducing  sulphuric 
acid.  When  the  acid  is  heated  with  certain  elements  which  have 
an  affinity  for  oxygen,  1  atom  of  the  latter  is  lost  and  the  sulphuric 
acid,  H2S04,  is  reduced  to  sulphurous  acid,  H2SO3,  which  in  turn 
breaks  down  into  sulphur  dioxide  and  water.  The  reaction  when 
copper  is  used  can  be  considered  to  take  place  according  to  the 
steps  represented  by  the  following  equations: 

H2SO4  =  H2SO8  +  O  (1) 

.       H2SO3  =  SO2  +  H2O  (2) 

Cu  +  O  =  CuO  (3) 

CuO  +  H2SO4  =  CuS04  +  H20  (4) 


258  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  reactions  thus  indicated  take  place  simultaneously.  Oxygen 
is  not  set  free  as  such,  but  unites  with  the  copper  to  form  copper 
oxide,  which,  being  the  oxide  of  a  metal,  reacts  in  turn  with  some 
of  the  sulphuric  acid  to  form  copper  sulphate.  It  will  be  recalled 
that  if  oxygen  gas  is  produced  by  a  chemical  reaction  the  formula 
used  for  it  in  the  equation  expressing  the  reaction  is  O2.  In  the  first 
and  third  equations  given  above  the  oxygen  is  represented  by  the 
symbol  O  in  order  to  indicate  that  these  equations  represent  a 
process  of  oxidation  and  that  none  of  the  gas  is  set  free.  All 
processes  of  oxidation  may  be  considered  as  taking  place  in  two 
steps — first,  the  breaking  down  of  the  oxidizing  agent  to  furnish 
oxygen,  and,  second,  the  oxidation  of  the  element  or  compound 
present.  The  four  equations  represent  a  single  process  and  may 
be  combined  into  one.  The  method  by  which  this  is  done  will  be 
described  fully  as  many  reactions  will  be  met  with  which  are  best 
considered  as  the  result  of  the  summation  of  several  changes  taking 
place  at  the  same  time. 

In  a  series  of  partial  equations  like  those  under  discussion,  if  a 
symbol  or  formula  appears  on  one  side  of  one  equation  and  on  the 
other  side  of  another,  it  indicates  that  the  substance  is  first  formed 
and  then  reacts.  As  a  consequence,  it  is  not  present  in  the  final 
product  and  its  formula  does  not  appear  in  the  completed  equa- 
tion for  the  reaction.  In  the  case  given  above,  oxygen  for  the 
reaction  is  furnished  by  the  sulphuric  acid;  and  it  is  used  up  by 
combining  with  the  copper.  We  are  justified,  then,  in  canceling 
the  O  on  the  right-hand  side  of  the  first  equation  and  the  O  on  the 
left-hand  side  of  the  third  equation.  For  the  same  reason  the 
H2SO3  in  equations  1  and  2,  and  the  CuO  in  equations  3  and  4  can 
be  canceled.  The  next  step  in  arriving  at  the  final  equation  is  to 
add  all  the  formulas  that  remain  on  the  left  side  of  the  partial 
equations  and  make  them  equal  the  sum  of  the  uncanceled  for- 
mulas on  the  right  side.  The  process  yields  the  following,  which 
is  the  equation  that  represents  the  reaction  between  copper  and 
concentrated  sulphuric  acid: 

Cu  +  2H2SO4  =  SO2  +  CuSO4  +  2H2O 

In  order  to  check  the  result,  the  final  equation  should  be  inspected 
to  see  if  only  the  formulas  of  the  substances  used  appear  on  one  side 
and  only  the  formulas  of  those  actually  obtained  appear  on  the 


THE  OXIDE  AND  ACIDS  OF  SULPHUR  259 

other.  The  number  of  atoms  of  each  element  appearing  on  the  two 
sides  should  also  be  noted  to  see  if  the  equation  balances. 

The  process  described  above  may  appear  to  be  a  long  and 
involved  one  to  arrive  at  such  a  simple  result.  It  is  probable 
that  the  student  could  have  written  the  final  equation  directly,  if 
he  had  known  what  were  the  products  formed  in  the  reaction.  The 
equations  used  as  an  example  in  this  case  are  very  simple  ones, 
but  others  will  be  met  with  later  which  are  more  complicated  and 
can  be  interpreted  best  by  the  use  of  the  method  described. 

288.  Sulphur  dioxide  is  formed  when  concentrated  sulphuric 
acid  is  heated  with  a  number  of  the  less  active  metals,  for  example, 
lead  or  mercury.  In  the  case  of  the  more  active  metals  the  acid  is 
reduced  farther.  Copper  removes  but  1  oxygen  atom,  as  we  have 
seen,  and  H2SO3(SO2  -f  H2O)  is  formed.  When  zinc  reacts  with 
it,  however,  all  the  oxygen  is  removed  and  hydrogen  sulphide  is 
produced.  Sulphuric  acid  does  not  oxidize  the  least  active  metals, 
such  as  platinum  and  gold.  For  this  reason  pans  of  platinum  are 
used  when  dilute  sulphuric  acid  is  concentrated  by  boiling  off  the 
water  which  it  contains. 

Sulphur  dioxide  is  formed  when  a  number  of  non-metallic 
elements  are  heated  with  concentrated  sulphuric  acid.  With 
sulphur  the  reaction  is  expressed  by  the  following  equations: 

H2SO4  =  H2S03  +  O 
H2SO3  =  SO2  +  H2O 

S  +  2O  =  S02 

There  is  not  a  fourth  partial  equation  in  this  case,  as  there  is  when 
copper  is  used,  because  sulphur  dioxide  is  not  an  oxide  of  a  metal 
and  does  not,  therefore,  react  with  the  sulphuric  acid  present  to 
form  a  sulphate.  Before  the  equations  are  combined  the  first  one 
must  be  multiplied  by  2,  since  2  oxygen  atoms  are  required  for  the 
oxidation  of  a  sulphur  atom,  according  to  the  third  equation.  Since 
no  oxygen  is  set  free  there  must  be  the  same  number  of  atoms  of 
the  element  on  the  two  sides  of  the  equations.  As  the  result  of 
this  change  two  molecules  of  sulphurous  acid,  H2SO3,  are  produced. 
This  fact  makes  it  necessary  to  multiply  the  second  equation  by  2, 
since  all  the  sulphurous  acid  breaks  down  into  sulphur  dioxide  and 
water.  The  equations  can  now  be  rewritten.  It  adds  to  the  clear- 


260  INORGANIC  CHEMISTRY  FOR  COLLEGES 

ness  to  enclose  in  brackets  the  formulas  of  the  substances  which  do 
not  appear  in  the  final  equation.  With  these  changes  the  equa- 
tions are  as  follows : 

2H2SO4  -  [2H2S03]  +  [2O] 
[2H2SO3]  =  2SO2  +  2H2O 
S  +  [2O]  =  SO2 

The  addition  of  these  partial  equations  leads  to  the  following 
equation,  which  expresses  the  reaction  between  sulphur  and  con- 
centrated sulphuric  acid: 

2H2SO4  +  S '  =  3SO2  +  2H2O 

289.  Physical  Properties  of  Sulphur  Dioxide.— Sulphur  dioxide 
is  a  colorless  gas,  which  is  slightly  more  than  twice  as  heavy  as  air. 
It  can  be  readily  recognized  by  its  odor,  which  is  that  produced 
when  sulphur  burns.     The  gas  can  be  condensed  to  a  colorless 
liquid  by  pressure  alone,  since  its  critical  temperature  is  156°. 
The  liquid  freezes  at  —73°  and  boils  at  —10°;  it  can  be  readily 
obtained  from  the  gas  by  passing  the  latter  through  a  vessel  sur- 
rounded by  a  mixture  of  ice  and  salt.     Liquid  sulphur  dioxide  is 
put  on  the  market  in  stout  iron  cans.     This  is  possible  because 
the  pressure  of  the  gas  generated  from  the  liquid  at  ordinary 
temperatures  is  not  very  great;  at  20°  it  is  3.24  atmospheres. 
The  gas  is  relatively  soluble  in  water;  50  volumes  under  ordinary 
conditions  dissolve  in  1  volume  of  water. 

290.  Chemical  Conduct  of  Sulphur  Dioxide. — The  heat  pro- 
duced when  solid  sulphur  burns  is  71,000  calories  for  each  gram- 
atom  (32  grams)  of  sulphur.     As  this  number  is  relatively  large 
we  should  expect  sulphur  dioxide,  the  product  of  the  reaction,  to  be 
stable   toward   heat.      It   dissociates   into   its  constituents   only 
slightly  at  high  temperatures,  and  in  this  behavior  resembles  water 
and  other  stable  compounds.    The  chief  chemical  properties  of 
sulphur  dioxide  are  due  to  the  fact  that  sulphur  in  many  of  its 
compounds  has  a  valence  of  6;   in  sulphur  dioxide  the  valence  of 
the  element  is  4,  and,  as  a  result,  when  the  compound  is  brought 
into  contact  with  certain  other  substances  chemical  reaction  takes 
place.     With  chlorine,  for  example,  sulphuryl  chloride,  SO2C12,  is 
formed  by  direct  combination,  the  reaction  being  facilitated   by 


THE  OXIDES  AND  ACIDS  OF  SULPHUR  261 

the  presence  of  a  trace  of  camphor,  which  serves  as  a  catalytic 
agent.  Sulphuryl  chloride  is  a  liquid  which  boils  at  69°;  it  decom- 
poses rapidly  with  water  to  form  sulphuric  and  hydrochloric  acids. 

Sulphur  dioxide  enters  into  a  number  of  reactions  of  this  type, 
the  most  important  of  which  is  the  one  that  takes  place  between 
the  gas  and  oxygen.  As  a  result  of  the  union  sulphur  trioxide, 
SOs,  is  formed.  The  reaction  is  one  of  great  interest  and  will  be 
discussed  fully  later,  as  it  is  the  basis  of  one  of  the  most  important 
of  the  chemical  industries — the  manufacture  of  sulphuric  acid. 

Sulphur  dioxide  unites  with  water  to  form  sulphurous  acid: 

SO2  +  H2O  <±  H2SO3 

The  reaction  indicated  by  the  equation  above  is  a  reversible  one; 
in  an  aqueous  solution  of  sulphur  dioxide  there  is  always  present 
some  of  the  free  oxide  and  some  of  the  acid.  It  has  been  found 
impossible  to  isolate  the  latter  on  account  of  the  fact  that  it  decom- 
poses when  an  attempt  is  made  to  separate  it  from  solution.  The 
sulphurous  acid  present  in  the  solution  takes  up  oxygen  from  the 
air  and  from  certain  compounds  that  contain  the  element,  and 
sulphuric  acid,  H2SO4,  is  formed.  This  reaction  of  sulphur  dioxide, 
in  which  it  is  a  reducing  agent,  is  made  use  of  in  the  process  of 
bleaching,  which  will  be  described  later. 

When  brought  into  contact  with  certain  substances  sulphur 
dioxide  loses  its  oxygen;  with  these  it  acts  as  an  oxidizing  agent. 
Only  the  most  powerful  reducing  agents,  however,  can  effect  the 
reduction.  One  case  has  already  been  noted — the  reaction  between 
sulphur  dioxide  and  hydrogen  sulphide : 

2H2S  -f  SO2  =  3S  +  2H2O 

In  this  case  the  hydrogen  sulphide  is  oxidized  and  the  sulphur  diox- 
ide reduced. 

Hydriodic  acid  is  a  powerful  reducing  agent  which  removes 
oxygen  from  sulphur  dioxide: 

4HI  +  SO2  =  2H2O  +  S  +  2I2 

291.  Uses  of  Sulphur  Dioxide. — Large  quantities  of  the  gas  are 
used  in  the  manufacture  of  sulphuric  acid,  which  is  of  fundamental 
importance  in  many  chemical  industries. 


262  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Sulphur  dioxide  finds  another  important  use  in  the  bleaching 
of  straw,  wool,  silk,  and  other  substances  which  are  affected  dele- 
teriously  by  chlorine.  The  chemistry  of  the  process  has  already 
been  briefly  described.  In  the  presence  of  moisture  the  sulphurous 
acid  first  formed  withdraws  oxygen  from  the  coloring  matters 
present;  as  a  result  of  the  reduction  colorless  compounds  are 
formed,  and  the  product  is  said  to  be  bleached.  It  is  probable 
that  in  the  case  of  straw  and  certain  other  substances,  the  bleaching 
is  produced  as  the  result  of  the  union  of  sulphur  dioxide  with  the 
colored  compounds  present.  The  resulting  product  is  colorless 
but  after  exposure  to  light  and  the  weather  for  some  time,  it 
undergoes  decomposition  and  the  original  coloring  matter  is 
reformed.  For  this  reason  straw  hats,  which  are  bleached  by 
exposure  in  a  closed  room  to  sulphur  dioxide  from  burning  sulphur, 
become  yellow  with  age. 

When  wool  is  bleached  it  is  first  scoured  to  remove  adhering 
grease  and  then,  while  still  moist,  it  is  subjected  to  the  action  of 
sulphur  dioxide  produced  by  burning  sulphur.  A  more  convenient 
process,  which  is  often  used,  is  to  pass  the  wool  through  a  dilute 
solution  of  a  sodium  bisulphite  and  then  through  one  of  sulphuric 
acid.  The  acid  decomposes  the  salt  and  liberates  sulphur  dioxide 
throughout  the  fiber. 

Sulphur  dioxide  was  formerly  much  used  as  a  disinfectant,  as 
it  kills  the  germs  of  disease.  The  gas  was  produced  by  burning 
sulphur  in  the  room  to  be  disinfected  after  windows  and  doors  had 
been  sealed.  As  the  gas  often  bleached  colored  materials  in  the 
room,  and  more  or  less  sulphuric  acid  was  formed,  the  process  was 
given  up  and  formaldehyde  was  substituted  for  the  purpose. 

Sulphur  dioxide  is  used  as  a  preservative,  as  it  prevents  the 
growth  of  bacteria.  For  this  purpose  it  has  been  used  in  wines 
and  beer  and  in  preserving  fruit  which  had  to  be  shipped  long  dis- 
tances. Even  small  quantities  of  sulphur  dioxide  are  said  to  be 
deleterious  to  health  and  the  practice  is  now  prohibited  by  law. 

Large  quantities  of  sulphur  dioxide  are  used  in  making  salts 
of  sulphurous  acid.  The  most  important  of  these  are  the  acid 
sodium  salt,  NaHSOs,  which  is  a  source  of  sulphur  dioxide  for 
bleaching,  as  has  been  mentioned,  and  the  corresponding  calcium 
salt  which  is  used  in  converting  wood  into  pulp  for  the  manufac- 
ture of  paper. 


THE  OXIDES  AND  ACIDS  OF  SULPHUR  263 

292.  Sulphurous  Acid. — It  has  already  been  stated  that  sul- 
phur dioxide  when  dissolved  in  water  reacts  with  the  solvent  and  is, 
in  part,  converted  into  sulphurous  acid.    The  solution  shows  the 
characteristic  properties  of  the  solutions  of  all  acids;   it  has  a  sour 
taste,  turns  blue  litmus  red,  reacts  with  metals,  and  forms  salts 
with  bases.     The  chemical  conduct  of  sulphur  dioxide  in  this 
respect  recalls  that  of  carbon  dioxide.     Like  the  carbonates,  the 
salts  of  sulphurous  acid — the  sulphites — are  decomposed  by  other 
acids  and  the  anhydride  of  the  acid  (SO2)  is  liberated.     Equa- 
tions for  two  typical  reactions  which  illustrate  the  analogy  are  as 
follows : 

Na2C03  +  2HC1  =  2NaCl  +  CO2  +  H2O 

Na2S03  +  2HC1  =  2NaCl  +  SO2  +  H2O 

The  explanation  of  the  reaction  in  the  case  of  sodium  sulphite  is 
similar  to  that  offered  for  the  formation  of  carbon  dioxide  from  a 
carbonate  (200);  sulphurous  acid  is  first  formed  as  the  result  of 
double  decomposition  and  then  breaks  down  into  sulphur  dioxide 
and  water. 

Sulphurous  acid  is  a  dibasic  acid;  the  two  hydrogen  atoms 
present  in  it  can  be  replaced  by  metallic  atoms.  The  salts  of  the 
acid  are  called  sulphites  in  accordance  with  the  system  of  nomen- 
clature adopted  by  chemists.  If  the  name  of  the  acid  ends  in  the 
syllable  ous,  the  name  of  the  salt  ends  in  the  syllable  ite.  Thus, 
sulphurous  acid  furnishes  sulphites,  chlorous  acid  chlorites,  and 
arsenous  acid  arsenites. 

293.  Chemical  Conduct  of  Sulphurous  Acid.— The  solution  of 
sulphur  dioxide  in  water,  which  contains  sulphurous  acid,  is  an 
active  reducing  agent.     The  reaction  which  takes  place  between 
it  and  chlorine  is  typical: 

C12  +  H2SO3  +  H2O  =  H2SO4  +  2HC1 

In  this  case  the  sulphurous  acid  has  been  oxidized  to  sulphuric  acid, 
and  the  chlorine  reduced  to  hydrochloric  acid.  The  salts  of  sul- 
phurous acid  behave  in  a  similar  way;  they  are  used  to  destroy 
the  excess  of  chlorine  when  the  latter  is  employed  as  a  bleaching 
agent. 

294.  Sulphites. — There  are  two  types  of  sulphites — acid  salts 
in  which  but  1  hydrogen  atom  of  sulphurous  acid,  H2S03,  has  been 


264  INORGANIC  CHEMISTRY  FOR  COLLEGES 

replaced  by  a  metallic  atom,  and  neutral  or  normal  salts  in  which 
both  hydrogen  atoms  have  been  replaced.  The  method  of  naming 
these  salts  is  the  same  as  that  used  in  the  case  of  carbonic  acid. 
The  salt  of  the  formula  NaHSOs  is  called  sodium  hydrogen  sul- 
phite, acid  sodium  sulphite,  or  sodium  bisulphite. 

A  number  of  sulphites  are  used  commercially.  Reference  has 
been  made  to  the  sodium  and  calcium  salts.  The  normal  so- 
dium salt,  Na2SOs,  either  anhydrous  or  in  the  hydrated  form, 
Na2SO3,5H2O,  is  used  in  photography  as  a  constituent  of  develop- 
ing solutions  for  films  or  paper.  These  solutions  are  powerful 
reducing  agents  and  take  up  oxygen  rapidly  from  the  air.  If  this 
is  permitted  they  soon  lose  their  power  as  developers  and  turn 
dark  and  stain  the  film.  If  a  sulphite  is  present  this  is  largely 
prevented,  because  it  absorbs  the  oxygen,  being  changed  into  a 
sulphate,  and  thus  protects  the  developing  agent  from  oxidation. 

Potassium  metabisulphite  is  often  used  in  photographic  work 
instead  of  the  neutral  sulphite.  It  has  the  formula  K2S2O.5,  and 
may  be  considered  as  formed  from  2  molecules  of  potassium 
bisulphite  as  the  result  of  the  loss  of  1  molecule  of  water: 

2KHS03  =  K2S205  +  H2O 

The  prefix  meta  is  often  used  in  inorganic  chemistry  in  naming 
acids  produced  from  other  acids  as  the  result  of  the  loss  of  water; 
for  example,  HaPCU  is  the  formula  of  phosphoric  acid  and  HPOs 
that  of  metaphosphoric  acid;  the  latter  is  formed  as  the  result  of 
the  loss  of  a  molecule  of  water  from  the  former.  Potassium 
metabisulphite  crystallizes  well,  is  relatively  stable  when  dry,  and 
weight  for  weight  absorbs  more  oxygen  than  sodium  sulphite. 

The  solubility  of  the  normal  sulphites  of  the  common  metals 
resembles  that  of  the  carbonates.  All  are  insoluble  or  difficultly 
soluble  in  water  except  those  of  sodium,  potassium,  and  ammonium. 

The  action  of  heat  on  most  sulphites  brings  about  a  decomposi- 
tion into  sulphur  dioxide  and  the  oxide  of  the  metal  present.  The 
temperature  at  which  this  takes  place  varies  with  the  metal.  As  is 
the  case  with  the  carbonates  the  sulphites  of  the  less  active  metals 
are  most  readily  decomposed.  The  salts  of  sodium  and  potassium 
decompose  at  very  high  temperatures  only  and  in  these  cases  a  mix- 
ture of  sulphide  and  sulphate  is  formed.  If  the  heating  is  done  in 
the  presence  of  air,  sulphites  take  on  oxygen  and  are  converted 


THE  OXIDES  AND  ACIDS  OF  SULPHUR  265 

into  sulphates,  which,  in  general  decompose  at  a  higher  tempera- 
ture than  sulphites. 

The  test  for  sulphites  is  carried  out  by  treating  the  compound 
with  dilute  hydrochloric  acid  and  noting  the  odor  of  the  evolved 
gas. 

SULPHUR  TRIOXIDE 

295.  Preparation  of  Sulphur  Trioxide. — When  sulphur  burns 
in  the  air,  the  product  is  sulphur  dioxide,  but  it  is  highly  probable 
that  more  or  less  of  the  trioxide  is  also  formed.     At  the  high  tem- 
perature of  the  flame  the  latter  decomposes  into  the  dioxide  and 
oxygen.     In  confirmation  of  this  view  it  has  been  shown  that  if  the 
air  in  which  the  sulphur  burns  contains  water-vapor,  appreciable 
quantities  of  the  trioxide  can  be  detected  as  the  result  of  the 
reaction. 

The  amount  of  sulphur  trioxide  formed  when  sulphur  burns  is 
small,  and  the  reaction  does  not  serve,  therefore,  as  a  means  of 
preparing  it.  Sulphur  trioxide  is  formed  when  the  sulphates  of 
certain  metals  are  heated.  The  substance  was  formerly  prepared 
in  this  way,  but  it  is  now  produced  directly  as  the  result  of  the 
union  of  sulphur  dioxide  and  oxygen  of  the  air  in  the  presence  of 
platinum.  As  the  preparation  of  the  oxide  in  this  way  is  the  most 
important  step  in  one  of  the  methods  used  to  manufacture  sul- 
phuric acid,  it  will  be  described  in  detail  later. 

296.  Physical  Properties  of  Sulphur  Trioxide. — There  are  two 
forms  of  sulphur  trioxide;    the  one  obtained  as  the  result  of  the 
union  of  sulphur  dioxide  and  oxygen  in  the  presence  of  finely 
divided  platinum  and  in  the  absence  of  water,  is  a  colorless  liquid 
which  boils  at  46°  and  freezes  at  15°  to  a  solid  resembling  glass  in 
appearance.     If  a  trace  of  moisture  is  added,  some  sulphuric  acid 
is  formed,  and,  as  a  result  of  the  presence  of  a  small  amount  of  acid, 
the  liquid  changes  on  standing  to  a  mass  of  white,  opaque,  needle- 
shaped  crystals.     These  do  not  melt  when  heated,  but  pass  into  a 
vapor,  which  on  condensation  yields  the  liquid  variety  of  the  tri- 
oxide.    The  molecular  weights  of  the  two  forms  have  been  deter- 
mined by  the  freezing-point  method;   from  these  it  has  been  cal- 
culated that  the  formula  of  the  liquid  variety  is  SOs  and  that 
of  the  solid  form  82(^6.     The  white  crystalline  variety  of  sulphur 
trioxide  is  said  to  be  a  polymer  of  the  liquid  form.     It  should  be 


266  INORGANIC  CHEMISTRY  FOR  COLLEGES 

called  sulphur  hexoxide,  as  it  is  really  a  different  substance  from 
sulphur  trioxide.  Since  it  passes  so  readily  into  the  trioxide  and  its 
chemical  behavior  is  identical  with  that  of  the  latter,  the  dif- 
ference is  not  emphasized  in  its  name. 

297.  Chemical  Behavior  of  Sulphur  Trioxide. — The  combining 
power  of  an  element,  as  indicated  by  its  valence,  is  one  of  its  most 
important  chemical  characteristics.  It  determines  the  composi- 
tion of  the  compounds  derived  from  the  element,  and,  in  certain 
cases,  even  its  fundamental  chemical  behavior;  some  elements 
have  acid-forming  properties  when  they  exhibit  one  valence,  and 
base-forming  properties  when  they  show  another.  The  processes 
of  oxidation  and  reduction  are  associated  with  change  in  valence, 
and  many  other  important  facts  of  chemical  interest  can  be  inter- 
preted through  a  study  of  the  valencies  of  the  elements  involved. 

A  comparison  of  the  behavior  of  sulphur  dioxide  and  sulphur 
trioxide  when  heated  to  high  temperatures  brings  out  a  conclusion 
which  can  be  applied  to  other  compounds.  The  dioxide  is  very 
stable  towards  heat;  the  trioxide,  on  the  other  hand,  begins  to 
decompose  below  400°  into  the  dioxide  and  oxygen,  and  as  the  tem- 
perature is  raised  the  decomposition  rapidly  increases.  The 
reaction  which  takes  place  is  a  reversible  one, 

2SO3  ^±  2SO2  +  O2 

and  the  equilibrium  is  markedly  affected  by  the  temperature. 
The  facts  in  regard  to  this  reaction  have  been  carefully  studied, 
since  it  is  the  basis  for  the  manufacture  of  sulphuric  acid, 
as  has  been  repeatedly  mentioned.  The  particular  fact  to  be 
emphasized  at  this  point  is  that  the  combining  power  of  an  ele- 
ment is  markedly  affected  by  temperature;  with  rise  in  tempera- 
ture the  ability  of  an  element  to  hold  others  in  combination  with 
it,  falls  off.  This  is  a  generalization  which  applies  to  all  elements, 
and  is,  therefore,  of  prime  importance. 

Sulphur  trioxide  reacts  with  water  violently;  when  a  small 
amount  of  it  is  dropped  into  the  liquid  a  hissing  sound  results, 
like  that  produced  when  red-hot  iron  is  thrust  into  water.  Since 
sulphuric  acid  is  the  product  of  the  reaction  sulphur  trioxide  is 
often  called  sulphuric  anhydride: 

SO3  +  H20  =  H2S04 


THE  OXIDES  AND  ACIDS  OF  SULPHUR  267 

On  account  of  its  great  affinity  for  water  sulphur  trioxide  fumes 
in  the  air  (141).  When  the  vapor  comes  into  contact  with  moist 
air  a  cloud  of  great  density  is  formed.  This  fact  was  utilized  in 
the  recent  war  to  produce  smoke  screens  for  use  along  the  battle 
front  and  to  protect  vessels  at  sea.  To  produce  a  dense  cloud  on 
land  sulphuric  acid  containing  sulphur  trioxide  in  solution,  so 
called  fuming  sulphuric  acid,  was  allowed  to  run  on  lime.  The 
reaction  which  took  place  between  the  acid  and  the  base  produced 
sufficient  heat  to  volatilize  the  sulphur  trioxide,  which  produced  a 
dense  white  smoke  with  the  moisture  in  the  air.  To  make  a 
smoke  screen  at  sea  the  fuming  acid  was  introduced  into  the 
smoke-stack  of  the  vessel  where  the  hot  gases  from  the  fires  under 
the  boiler  vaporized  the  sulphur  trioxide. 

SULPHURIC  ACID 

298.  Sulphuric  acid  is  a  fundamental  raw-material  in  many  of 
the  largest  chemical  industries,  and  in  others  in  which  it  is  not 
itself  used,  substances  are  employed  that  require  it  in  their  prep- 
aration. Sulphuric  acid  is  thus  the  most  widely  used  manufac- 
tured compound  in  industrial  chemistry.  Liebig,  one  of  the  leaders 
in  the  early  days  of  chemistry,  said  the  civilization  of  a  nation 
could  be  gauged  by  the  amount  of  soap  it  consumed  per  capita; 
later  it  was  claimed  that  sulphuric  acid  should  be  the  standard  of 
measure,  for  we  can  trace  back  to  it  so  many  of  the  essentials  and 
conveniences  of  material  existence.  The  manufacture  of  the  acid 
has  been  one  of  the  great  problems  of  industrial  chemistry,  and 
many  chemists  have  exercised  their  skill  and  ingenuity  in  the  solu- 
tion of  the  problem,  which  is  a  very  complicated  one. 

Sulphuric  acid  was  known  in  ancient  times  because  it  is  pro- 
duced by  the  action  of  water  on  the  sulphur  trioxide  formed  when 
many  naturally  occurring  sulphates  are  heated  to  a  high  tempera- 
ture. Fuming  sulphuric  acid  (H2SO4  +  SOs)  was  manufactured 
in  the  middle  of  the  eighteenth  century  at  Nordhausen,  in  Ger- 
many, by  heating  ferrous  sulphate,  FeSO4,  7H2O,  but  the  product 
was  very  expensive.  The  common  name  of  the  salt  was  iron 
vitriol,  and  the  oily  product  produced  from  it  was  called  oil  of 
vitriol — a  name  still  used  in  chemical  trade.  In  1740  Ward,  in 
England,  started  the  manufacture  of  the  acid  by  burning  sulphur 


268  INORGANIC  CHEMISTRY  FOR  COLLEGES 

with  niter,  KNOs,  in  the  presence  of  water  under  great  glass 
domes.  A  few  years  later  lead  chambers  were  substituted  for  the 
glass  domes,  and  in  1793  it  was  shown  that  the  oxidation  of  sulphur 
dioxide  in  the  process  was  brought  about  through  the  catalytic 
influence  of  the  oxides  of  nitrogen  formed  from  the  nitrate  used. 

299.  Manufacture  of  Sulphuric  Acid. — The  preparation  of  the 
acid  involves,  in  the  main,  but  a  few  simple  reactions;  sulphur,  or  a 
sulphide,  is  burned  to  sulphur  dioxide,  which  is  then  oxidized,  and 
the  resulting  compound  is  treated  with  water.  It  has  already  been 
shown  how  this  can  be  done,  but  serious  difficulties  are  met  with 
when  the  process  is  carried  out  on  an  industrial  scale.  The  reac- 
tion as  the  result  of  which  sulphur  dioxide  unites  with  oxygen  and 
forms  sulphur  trioxide  in  the  presence  of  catalytic  agents,  has  been 
known  for  a  long  time,  but  it  was  only  in  1901  that  it  was  applied 
to  the  manufacture  of  sulphuric  acid.  As  the  result  of  extended 
chemical  research  a  method  was  worked  out  for  the  synthesis  of 
indigo — our  most  important  blue  dye.  This  method  involved,  in 
one  of  the  steps,  the  use  of  fuming  sulphuric  acid.  To  compete 
successfully  with  the  natural  dyestuff  obtained  from  a  plant  grown 
in  India,  it  was  necessary  to  reduce  the  cost  of  the  production  of 
sulphur  trioxide.  In  a  search  for  a  new  method,  the  direct  oxida- 
tion of  sulphur  dioxide  to  the  trioxide  in  the  presence  of  catalytic 
agents  was  carefully  studied,  and  as  the  result  the  so-called 
"  contact  "  process  for  the  manufacture  of  sulphuric  acid  was 
worked  out.  While  the  method  was  primarily  devised  for  the 
preparation  of  sulphur  trioxide  and  fuming  sulphuric  acid,  it  was 
readily  applied  to  the  manufacture  of  the  ordinary  acid.  The 
older,  more  cumbersome  process  is  still  in  use  to-day,  but  it  is 
probable  that  as  new  plants  are  erected  the  contact  process  will 
be  employed,  except  in  the  case  where  the  nature  of  the  raw 
materials  available  favor  the  use  of  the  older  method. 

The  oxidation  of  sulphur  dioxide  to  trioxide  takes  place  very 
slowly  under  ordinary  conditions  and  it  is  necessary  to  increase 
the  rate  of  the  reaction  when  it  is  used  in  a  manufacturing  opera- 
tion. In  the  contact  process  the  catalytic  agent  which  increases 
the  rate  is  platinum;  in  the  older,  so-called  "  chamber  "  process 
certain  oxides  of  nitrogen  are  used  for  this  purpose.  The  reactions 
which  take  place  in  the  latter  process  are  more  or  less  complicated, 
as  we  shall  see,  but  in  the  main  they  involve  the  oxidation  of  the 


THE  OXIDES  AND  ACIDS  OF  SULPHUR  269 

sulphur  dioxide  by  nitrogen  dioxide,  NC>2,  which  is  reduced,  as  a 
result,  to  nitric  oxide,  NO.  This  oxide  unites  directly  with  the 
oxygen  in  the  air  to  form  nitrogen  dioxide,  which,  in  turn,  oxidizes 
more  sulphur  dioxide.  The  nitric  oxide  serves  thus  as  a  carrier 
of  the  oxygen  from  the  air  to  the  sulphur  dioxide;  at  one  instant 
the  nitrogen  is  in  the  form  of  NO  and  at  the  next  of  NO2,  then 
NO,  then  NO2,  and  so  on.  The  word  carrier  has  been  selected 
to  name  this  process,  which,  it  will  be  seen,  is  a  catalytic  one,  for 
the  nitrogen  compounds  involved  do  not  enter  into  the  final 
product  of  the  reaction. 

300.  Contact  Process  for  Sulphuric  Acid. — The  equation  for  the 
reaction  upon  which  this  process  is  based  is  as  follows : 

2SO2  +  O2  «=*  2SO3  +  41,800  cal. 

The  equation  shows  that  the  reaction  is  reversible,  and  indicates 
the  amount  of  heat  evolved  when  it  takes  place.  Both  of  these 
facts  are  of  prime  importance.  To  get  the  best  result,  the  reac- 
tion must  be  carried  out  under  the  conditions  where  the  equilib- 
rium is  such  that  the  largest  practical  amount  of  sulphur  trioxide 
is  formed.  These  conditions  are  determined  by  the  temperature 
at  which  the  reaction  takes  place;  and  since  heat  is  evolved,  the 
proper  conditions  can  be  obtained  only  by  carrying  the  process 
out  in  such  a  way  that  the  heat  can  be  controlled.  Before  the 
reaction  can  be  adapted  to  the  commercial  preparation  of  sulphur 
trioxide,  the  effect  of  temperature  on  the  equilibrium  must  be 
known,  and  the  effect  of  various  catalyzing  agents  on  the  rate  at 
which  the  reaction  proceeds  must  be  determined.  At  the  tempera- 
ture at  which  the  reaction  proceeds  rapidly  enough  in  the  absence 
of  a  catalyzer  to  make  it  useful,  the  equilibrium  is  such  that  but  a 
small  amount  of  sulphur  trioxide  is  present  in  the  gases.  At  400° 
when  twice  the  theoretical  amount  of  oxygen  is  used,  between  98 
and  99  per  cent  of  the  dioxide  is  converted  into  trioxide;  at  720°, 
the  conversion  is  about  60  per  cent  complete;  and  at  900°  prac- 
tically no  sulphur  trioxide  is  formed.  It  is  evident  that  the  reac- 
tion should  be  carried  out  at  approximately  400°  if  possible. 

A  large  amount  of  heat  is  produced  as  a  result  of  the  reaction, 
and  as  the  temperature  would  rise  above  400°  if  the  heat  were  not 
controlled,  it  is  necessary  to  construct  the  apparatus  in  which  the 
reaction  is  carried  out  in  such  a  way  that  the  excess  heat  over  that 


270 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


required  to  maintain  the  proper  temperature  is  carried  away. 
This  is  done  by  using  this  excess  to  pre-heat  the  gases  to  the  tem- 
perature at  which  the  reaction  takes  place. 

A  number  of  substances  were  studied  in  the  investigation  of  the 
influence  of  catalytic  agents  on  the  oxidation  of  sulphur  dioxide, 
and  it  was  found  that  finely  divided  platinum  and  an  oxide  of  iron, 
Fe2Oa,  were  the  only  ones  that  were  of  practical  significance. 
Catalyzers  have  no  effect  on  the  final  equilibrium  attained  in  a 
reaction,  but  differ  from  one  another  in  their  influence  on  the  rate 
of  the  reaction.  The  oxidation  of  sulphur  dioxide  takes  place 
rapidly  at  400°  in  the  presence  of  platinum,  and  at  625°  when 
ferric  oxide,  Fe2C>3,  is  used.  Since  at  the  latter  temperature  only 
about  70  per  cent  of  the  dioxide  is  converted  into  trioxide,  it  is 
evident  that  platinum  is  preferred  as  a  catalytic  agent  notwith- 
standing its  high  cost.  In  some  plants  two  contact  chambers  are 
used;  the  first  contains  ferric  oxide  and  the  second  platinum. 

301.  A  diagrammatic  sketch  of  the  lay-out  of  a  plant  for  the 
manufacture  of  sulphuric  acid  by  the  contact  process,  a  so-called 

WashingTpwer  ContactTower     Absorption  Tower 

I      DryingTower  j  \H2SO4 

Sulfur  Burner                         /             Conc     }       (0=$. 
I  Dust  Chamber /  ...  .       H2S04  /  , Jfc 


Air 


FIG.  29. 


flow-sheet,  is  given  in  Fig.  29.  The  various  steps  in  the  process 
can  be  seen  by  an  examination  of  the  diagram. 

The  gases  issuing  from  the  contact  chamber  are  passed  through 
strong  sulphuric  acid  (about  97  per  cent),  which  absorbs  the  sul- 
phur trioxide. 

302.  Lead  Chamber  Process  for  Sulphuric  Acid. — This  process, 
which  has  been  the  standard  one  for  over  a  hundred  years  in  chem- 
ical industry,  derives  its  name  from  the  fact  that  the  sulphuric 


THE  OXIDES  AND  ACIDS  OF  SULPHUK 


271 


acid  is  formed  in  chambers  built  of  lead.  The  mixing  of  the  gases 
involved  takes  place  rather  slowly  in  the  presence  of  the  large 
amount  of  nitrogen  in  the  product  resulting  from  the  burning  of 
sulphur  in  the  air;  and,  as  a  consequence,  the  reaction  chambers 
must  be  large.  They  are  built  of  lead,  as  this  metal  resists 
satisfactorily  the  acid  formed  in  the  process. 

Fig.  30  shows  diagrammatically  the  principal  features  of  a 
sulphuric  acid  plant.     Sulphur  dioxide  is  first  produced  by  burning 


* 

<— *  -NOandNOz 
TowerAcid. 


Nffmsyl  5u 
Acid 


s  ^JaoQQ~r\ 

I        L-fstLff***          ^ft^S^ 

tjyr  c°°'er   r  oxn 


FIG.  30. 


sulphur  or  by  roasting  a  sulphide,  usually  pyrite,  FeS2,  in  the  pres- 
ence of  air: 

4FeS2  +  11O2  =  2Fe2O3  +  8SO2 

The  gases  from  the  burners  traverse  a  flue  in  order  to  allow  the 
dust  present  to  settlej  and  are  there  mixed  with  the  required  amount 
of  air,  which  furnishes  the  oxygen  for  the  subsequent  oxidation. 
In  the  flue  is  placed  a  pot  containing  sodium  nitrate  and  sulphuric 
acid,  which  furnish  to  the  gases  enough  nitric  acid  to  produce  the 
small  amount  of  the  oxides  of  nitrogen  equivalent  to  that  lost  in 
the  subsequent  oxidation.  The  gases  next  pass  through  the  Glover 
tower,  which  will  be  explained  later,  pick  up  the  oxides  of  nitrogen 
there,  and  then  go  into  the  lead  chambers,  where  they  come  into 
contact  with  water,  which  is  introduced  at  several  places  in  the 
form  of  steam  or  as  a  fine  spray.  As  the  gases  slowly  drift  through 


272  INORGANIC  CHEMISTRY  FOR  COLLEGES 

they  are  converted  into  sulphuric  acid  which  collects  as  a  liquid  on 
the  floor  of  the  chambers. 

The  large  amount  of  free  nitrogen  present  in  the  gases  carries 
along  with  it  the  oxides  of  nitrogen.  In  order  to  prevent  the  loss 
of  the  latter  the  gas  issuing  from  the  chambers  is  passed  through 
the  Gay-Lussac  tower;  this  contains  tiles  or  coke  over  which  a 
stream  of  concentrated  sulphuric  acid  runs  continuously.  The 
acid  absorbs  the  oxides  of  nitrogen,  and  on  reaching  the  bottom  of 
the  tower  is  pumped  to  the  top  of  the  Glover  tower.  It  is  allowed 
to  flow  down  through  this  tower,  which  is  filled  with  pieces  of 
broken  flint,  along  with  some  of  the  dilute  acid  obtained  from  the 
lead  chambers.  The  compound  formed  from  the  oxides  of  nitrogen 
and  sulphuric  acid  in  the  Gay-Lussac  tower  is  decomposed  by  the 
water  present  in  the  dilute  acid,  and  the  oxides  thus  set  free.  The 
hot  gases  from  the  pyrite  burners  in  passing  through  the  Glover 
tower  carry  along  with  them  the  oxides  of  nitrogen  into  the  lead 
chambers.  As  the  dilute  acid  falls  through  the  Glover  tower  the 
hot  gases  in  passing  through  it  carry  off  most  of  the  water  present, 
and,  as  a  result,  the  sulphuric  acid  which  flows  from  the  bottom  of 
the  tower  contains  but  a  small  amount  of  water.  This  acid  is 
pumped  to  the  top  of  the  Gay-Lussac  tower  and  serves  to  collect 
the  oxides  of  nitrogen  as  before.  The  final  product  is  the  acid 
which  collects  in  the  chambers— the  so-called  chamber  acid. 

303.  The  chemical  reactions  which  take  place  in  the  chamber  process  for 
the  manufacture  of  sulphuric  acid  have  been  very  fully  studied.  Several 
explanations  of  the  way  in  which  the  oxides  of  nitrogen  act  have  been  put 
forward.  While  the  part  played  by  the  oxides  in  effecting  the  oxidation  of 
sulphur  dioxide  is  essentially  that  already  indicated,  the  reactions  are  more 
complex  than  those  given.  When  an  insufficient  amount  of  water  is  present 
in  the  lead  chambers  during  the  oxidation,  a  crystalline  compound  is  formed 
on  the  walls  of  the  chambers.  This  substance  is  supposed  to  be  an  inter- 
mediate product  in  the  reaction;  it  has  the  composition  represented  by  the 
formula  H(NO)SO4  and  is  called  nitrosyl  sulphuric  acid.  The  formula  is 
written  as  it  is  to  show  the  relation  between  the  composition  of  the  substance 
and  that  of  sulphuric  acid,  H2SO4;  it  may  he  considered  as  formed  by  the 
replacement  of  one  hydrogen  atom  in  the  acid  by  the  NO  group.  As  this 
group  appears  in  a  number  of  compounds  it  has  been  given  a  special  name  — 
nitrosyl.  The  formulas  of  sulphuric  acid  and  nitrosyl  sulphuric  acid  are 
often  written  in  the  following  manner: 


O=Q-  O  -  H  O=Q-  O  -  H 

0=°-  0  -  H  0=°-  O  - 


THE  OXIDES  AND  ACIDS  OF  SULPHUR  273 

It  Is  seen  that  the  NO  group  replaces  1  hydrogen  atom. 

Nitrosyl  sulphuric  acid  can  be  prepared  in  a  number  of  ways;  for  example, 
it  is  produced  as  the  result  of  the  direct  addition  of  nitric  acid  and  sulphur 
dioxide: 

SO2-f  HNO3  =  H(NO)SO4 

It  is  formed  when  nitric  oxide  and  nitrogen  dioxide  are  passed  into  concen- 
trated sulphuric  acid, 

(1)  2H2S04  +  NO+  NO,  <±  2H(NO)SO4  +  H2O 

and  when  sulphur  dioxide  is  oxidized  by  nitrogen  dioxide  in  the  presence  of 
water: 

(2)  H2O  +  2SO2  +  3NO2  =  2H(NO)SO4  +  NO 

With  these  equations  the  chief  reactions  which  take  place  in  the  formation 
of  sulphuric  acid  can  be  interpreted.  In  the  chamber,  reaction  (2)  probably 
first  takes  place.  Reaction  (1)  is  reversible  and,  as  a  result,  in  the  presence 
of  water  the  nitrosyl  sulphuric  acid  formed  breaks  down  into  sulphuric  acid 
and  the  oxides  of  nitrogen.  The  air  present  converts  the  nitric  oxide,  NO, 
into  the  dioxide  which  serves  to  react  with  more  sulphur  dioxide.  It  was 
stated  that  nitrosyl  sulphuric  acid  —  the  so-called  chamber  crystals  —  sepa- 
rates only  when  there  is  a  deficiency  of  water  in  the  chamber;  Equation  (1) 
furnishes  a  reason  for  this. 

The  gases  swept  along  by  the  nitrogen  pass  through  the  Gay-Lussac 
tower.  Here  the  reaction  represented  by  Equation  (1),  reading  from  left 
to  right,  takes  place.  It  is  seen  from  this  why  concentrated  sulphuric  acid 
is  needed  at  this  stage.  The  product  from  this  tower,  which  contains  nitrosyl 
sulphuric  acid  in  solution,  is  sent  through  the  Glover  tower,  where  it  comes 
in  contact  with  diluted  sulphuric  acid.  The  water  present  decomposes  the 
nitrosyl  sulphuric  acid,  as  indicated  by  Equation  (1),  reading  from  right  to 
left;  and  the  oxides  of  nitrogen  liberated  are  carried  along  with  the  sulphur 
dioxide  and  oxygen  into  the  contact  chambers,  where  the  cycle  is  begun  once 
more. 

By  considering  Equations  (1)  and  (2),  along  with  the  equation  for  the 
oxidation  of  nitric  oxide  to  nitrogen  dioxide,  as  partial  equations  in  a  single 
transformation,  we  can  bring  out  clearly  the  fact  that  the  nitrogen  compounds 
are  catalytic  agents.  The  three  equations  written  in  this  way,  the  formulas 
of  the  compounds  appearing  on  both  sides  of  the  equation  being  enclosed  in 
brackets  as  usual,  are  as  follows: 

H20+  2S02  +  [3N02]  -  [2H(NO)S04]  +  [NO] 
[2H(NO)SO4]  +  H-O=  2H2SO4  +  [NO]-f-  [NOJ 
[2NO]  +  O2  =  [2NO2] 


2H2O  +  2SO2  +  O2  =  2H2SO4 


274  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  equations  differ  from  the  partial  equations  studied  before  because,  in 
this  case,  they  represent  reactions  that  can  be  independently  realized;  the 
oxidation  of  sulphur  dioxide  represented  by  the  combined  equation  can  be 
carried  out  in  the  separate  steps  indicated  by  the  partial  equations. 

These  partial  equations  show  clearly  how  nitrogen  dioxide  serves  as  a 
catalytic  agent  in  effecting  the  oxidation.  In  the  case  of  the  use  of  platinum 
a  satisfactory  explanation  is  not  so  evident;  various  views  have  been  put 
forward,  but  adequate  experimental  proof  for  any  one  is  lacking.  It  was 
thought  at  one  time  that  the  platinum  formed  an  unstable  compound  with 
oxygen,  and  that  this  oxidized  the  sulphur  dioxide  and  was,  as  a  result, 
reduced  to  platinum;  the  metal  united  with  more  oxygen  and  the  cycle  was 
repeated.  According  to  this  explanation  the  platinum  serves  as  a  carrier 
in  the  same  general  way  that  nitric  oxide  does.  A  later  explanation  is  the 
more  probable  one.  According  to  this,  the  finely  divided  metal  adsorbs 
large  amounts  of  the  gases  on  its  surface  and  under  these  conditions  reaction 
between  them  takes  place  more  rapidly.  This  view  serves  to  explain  the 
behavior  as  catalytic  agents  of  very  inert  substances  like  powdered  glass  and 
sand,  where  the  formation  of  an  unstable  chemical  compound  with  the  reacting 
substances  is  highly  improbable. 

304.  The    product    of    the    chamber    process — the    so-called 
chamber  acid — contains  about  65  per  cent  of  sulphuric  acid  and 
varies  in  specific  gravity  between  1.5  and  1.6.    In  this  form  the 
acid  is  Used  in  a  number  of  commercial  operations,  but  for  others, 
it  is  concentrated  by  evaporating  off  most  of  the  water.     This 
can  be  carried  out  in  vessels  constructed  of  lead  until  the  con- 
centration reaches  77  per  cent,  when  the  specific  gravity  is  1.7. 
The  acid  reacts  slowly  with  lead  below  this  temperature,  but 
the  lead  sulphate  formed  adheres  closely  to  the  metal  and  serves  as 
a  protective  coating.     Hot  sulphuric  acid  stronger  than  77  per  cent 
dissolves  lead  sulphate,  and,  as  a  consequence,  the  acid  cannot  be 
concentrated  beyond  this  point  in  vessels  made  of  lead.     The  final 
evaporation  is  carried  out  either  in  silica  pans  or  stills  made  of 
cast  iron;  when  the  concentration  reaches  about  94  per  cent  and 
the  acid  has  the  specific  gravity  1.84  the  evaporation  is  stopped. 
The  product  is  the  concentrated  sulphuric  acid  of  commerce. 

Sulphuric  acid  is  shipped  in  steel  drums  or  tank  cars,  or  in 
carboys,  which  are  large  glass  bottles  packed  in  straw  in  wooden 
boxes,  the  neck  of  the  bottle  being  exposed  so  that  the  acid  can  be 
conveniently  poured  out.  The  usual  form  of  carboy  holds  about 
200  pounds  of  sulphuric  acid. 

305.  Physical  Properties  of  Sulphuric  Acid. — The  pure,  anhy- 
drous acid,  sometimes  called  hydrogen  sulphate,  is  an  oily  liquid 


THE  OXIDES  AND  ACIDS  OF  SULPHUR  275 

which  has  the  specific  gravity  1.84  at  15°,  is  miscible  with  water 
in  all  proportions,  and  freezes  to  a  crystalline  solid  that  melts  at 
10.5°.  When  heated  to  its  boiling-point,  336°,  it  undergoes  partial 
decomposition  and  a  mixture  of  sulphuric  acid  and  water  containing 
98.33  per  cent  of  the  acid  is  obtained.  This  mixture  boils  at  338° 
and  is  the  product  obtained  when  the  acid  is  distilled. 

Commercial  sulphuric  acid  contains  a  number  of  impurities; 
among  these  are  lead  sulphate,  arsenic  oxide,  and  oxides  of  nitrogen. 
The  presence  of  lead  sulphate  can  be  shown  by  diluting  the  acid 
with  water;  it  is  precipitated  because  it  is  insoluble  in  dilute 
sulphuric  acid.  The  so-called  C.P.  (chemically  pure)  acid  con- 
tains but  very  small  amounts  of  impurities,  the  presence  of  which 
can  be  neglected  in  most  of  the  cases  in  which  the  acid  is  used. 

306.  Chemical  Behavior  of  Sulphuric  Acid. — Sulphuric  acid  is 
an  active  chemical  reagent.  It  shows  a  number  of  reactions  of 
addition.  It  unites  directly  with  water  with  the  evolution  of  a 
large  amount  of  heat.  For  this  reason  care  must  be  taken  when 
the  acid  is  diluted  with  water.  If  the  latter  is  poured  on  the  acid 
it  floats  on  the  surface  and  the  large  amount  of  heat  generated 
causes  the  formation  of  steam  so  rapidly  that  a  lively  sputtering 
results;  if,  on  the  other  hand,  the  acid  is  poured  into  the  water, 
it  sinks,  the  two  liquids  mix  rapidly,  and  there  is  no  chance  of 
an  accident. 

Several  hydrates  of  sulphuric  acid  have  been  isolated;  the 
one  having  the  formula  H^SO^B^O  is  a  crystalline  compound 
which  melts  at  8°;  at  very  low  temperatures  hydrates  having  the 
composition  represented  by  the  formulas  H2SO4,2H20  and 
H2SO4,3H2O  are  formed.  The  affinity  of  sulphuric  acid  for  water 
is  so  great  that  it  is  an  excellent  agent  for  removing  water-vapor 
from  gases  that  do  not  react  with  it.  To  effect  this  the  gas  is 
simply  bubbled  through  the  liquid. 

Sulphuric  acid  decomposes  certain  compounds  containing 
hydrogen  and  oxygen  and  withdraws  from  them  these  elements  in 
the  proportion  to  form  water.  When  the  concentrated  acid,  for 
example,  is  warmed  with  sugar,  C^H^On,  the  principal  reaction 
which  takes  place  is  the  liberation  of  carbon  as  a  black  mass,  and 
the  formation  of  water  which  unites  with  the  sulphuric  acid.  A 
similar  reaction  takes  place  when  the  acid  is  allowed  to  stay  in 
contact  with  wood. 


276  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Sulphuric  acid  reacts  with  sulphur  trioxide  and  forms  disul- 
phuric  acid, 

H2SO4  +  SO3  <=±  H2S2O7 

which  is  present  in  the  so-called  fuming  sulphuric  acid;  it  readily 
breaks  down  into  its  constituents  and  the  sulphur  trioxide  fumes 
when  it  comes  in  contact  with  moist  air.  Salts  of  the  acid,  how- 
ever, are  stable;  they  are  usually  called  pyrosulphates1  and  are 
formed  by  heating  acid  sulphates  to  a  high  temperature : 

2NaHSO4  +±  Na2S2O7  +  H2O 

Sulphuric  acid  unites  directly  with  normal  salts  of  the  acid 
and  forms  acid  salts : 

H2SO4  +  Na2SO4  -  2NaHSO4 

Sulphuric  acid  enters  into  reactions  of  double  decomposition 
with  salts  of  other  acids,  and  is  a  very  valuable  reagent  in  the 
preparation  of  acids  having  boiling-points  lower  than  that  of 
sulphuric  acid  itself;  since  this  is  relatively  high,  many  acids  can 
be  prepared  in  this  way.  The  preparation  of  hydrochloric  acid 
is  an  example  that  is  already  familiar. 

Sulphuric  acid  enters  into  reactions  of  oxidation.  Some  of 
these  have  been  discussed  in  the  consideration  of  the  preparation 
of  sulphur  dioxide  (287).  All  the  metals  except  the  so-called 
noble  metals,  among  which  are  platinum  and  gold,  can  be  oxidized 
by  sulphuric  acid.  In  all  cases  sulphates  are  formed  as  the  final 
result  of  the  action  since  if  oxides  were  first  produced  they  would 
react  with  the  acid  present  to  form  salts.  The  product  formed  as 
the  result  of  the  reduction  of  sulphuric  acid  varies  with  the  activity 
of  the  metal  with  which  it  interacts.  The  more  active  metals 
like  zinc  and  magnesium  reduce  the  acid  to  hydrogen  sulphide: 

H2S04  =  H2S  +  [4O] 
4Zn  +  [4O]  =  [4ZnO] 
[4ZnO]  +  4H2SO4  =  4ZnSO4  +  4H2O 

4Zn  +  5H2SO4  =  H2S  +  4ZnSO4  +  4H2O 

1  The  prefix  pyro  is  derived  from  the  Greek  word  signifying  fire.  It  has 
been  used  often  in  naming  compounds  produced  as  the  result  of  the  action 
of  heat. 


THE  OXIDES  AND  ACIDS  OF  SULPHUR  277 

The  reaction  which  takes  place,  however,  is  more  complicated  than 
the  equation  indicates,  for  some  of  the  hydrogen  sulphide  formed  is 
oxidized  by  the  sulphuric  acid  and  sulphur  is  formed.  Hydrogen 
reduces  the  acid  to  sulphur  dioxide  at  about  160°. 

Sulphuric  acid  oxidizes  many  of  the  non-metals.  The  reaction 
with  sulphur  has  been  given  (288).  When  the  acid  is  heated  with 
carbon,  carbon  dioxide  is  formed: 

2H2SO4  =  2H2O  +  2SO2  +  [2O] 
C  +  [20]  -  C02 


2H2SO4  +  C  =  2H2O  +  2SO2  +  CO2 

The  activity  of  sulphuric  acid  as  an  oxidizing  agent  increases 
with  rise  in  temperature,  and  decreases  as  the  acid  is  diluted  with 
water. 

When  sulphuric  acid  is  dissolved  in  water  it  undergoes  ioniza- 
tion :  , 

H2S04  *=*  H+  +  HSO4~ 

HS04-  +*  H+  +  S04- 

The  extent  to  which  the  acid  is  ionized  is  determined  by  the  con- 
centration— that  is,  by  the  relative  amount  of  water  present. 
In  a  one-tenth  normal  solution  (4.9  grams  in  a  liter)  at  18°,  about 
61  per  cent  of  the  acid  is  converted  into  hydrogen  and  sulphate, 
S04~~,  ions.  Salts  of  sulphuric  acid  form  ions  in  a  similar  way; 
copper  sulphate  yields,  for  example,  a  copper  ion,  Cu++,  and  a 
sulphate  ion,  S04~~. 

307.  Sulphates. — Normal  sulphates  of  most  of  the  metals  are 
known,  and  many  of  them  have  important  industrial  applications, 
which  will  be  considered  later  in  connection  with  a  study  of  the 
metals  that  they  contain.  The  sulphates  which  occur  in  nature  as 
minerals  have  a  commercial  significance  and  are  mined  in  large 
quantities.  The  insoluble  sulphates  can  be  readily  formed  by 
double  decomposition  from  aqueous  solutions  of  salts,  one  of  which 
contains  the  sulphate  ion  and  the  other  the  ion  of  a  metal.  The 
soluble  sulphates  are  prepared  by  the  action  of  sulphuric  acid  on 
oxides,  hydroxides,  or  salts  of  acids  which  are  more  volatile  than 
sulphuric  acid,  for  example,  carbonates  and  chlorides. 


278  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  important  acid  sulphates  are  those  of  the  so-called  alkali 
metals,  sodium  and  potassium.  They  can  be  formed  by  the  par- 
tial neutralization  of  sulphuric  acid;  sodium  hydrogen  sulphate 
can  be  made  in  this  way: 

NaOH  +  H2SO4  =  NaHSO4  +  H2O 

Acid  salts  are  also  formed  by  treating  the  normal  salts  with  sul- 
phuric acid: 

Na2SO4  +  H2SOfi  =  2NaHS04 

These  salts  show  an  acid  reaction  in  solution,  since  they  yield 
hydrogen  ions  on  dissociation. 

The  sulphates  of  the  heavy  metals  decompose  at  high  tempera- 
tures into  sulphur  trioxide  and  the  oxide  of  the  metal  present,  for 
example, 

ZnSO4  =  ZnO  +  SO3 

The  sulphates  of  sodium  and  potassium  do  not  undergo  a  similar 
decomposition.  Acid  sulphates  on  heating  lose  water  and  pass 
into  pyrosulphates: 

2NaHSO4  =  Na2S2O7  +  H2O 

308.  Test  for  Sulphates.—  This  test  is  based  upon  the  fact  that 
barium  sulphate  is  the  only  common  barium  salt  that  is  insoluble 
in  dilute  nitric  acid.  When  a  solution  of  a  barium  salt  is  added  to 
a  solution  of  sulphuric  acid  or  a  sulphate;  a  white  precipitate  of 
barium  sulphate  —  a  very  insoluble  substance  —  is  formed  as  a 
result  of  double  decomposition  : 

Na2SO4  +  BaCl2  =  2NaCl  +  BaSO4 

When  this  equation  is  written  in  the  ionic  form,  as  follows, 
2Na+  +  SO4--  +  Ba++  +  2C1"  =  2Na+  +  BaSO4  +  2C1~ 


it  is  seen  that  the  reaction  consists  in  the  union  of  the  barium  and 
sulphate  ions  to  form  insoluble  barium  sulphate.  The  sodium 
chloride  formed  at  the  same  time,  being  soluble,  remains  in  solu- 
tion and  is,  consequently,  in  the  form  of  ions.  The  test  is,  there- 
fore, based  on  the  fact  that  if  a  barium  ion  is  brought  into  contact 
with  a  sulphate  ion,  the  two  unite  and  form  barium  sulphate, 
because  this  salt  is  insoluble.  There  are  many  barium  salts  whi<?h 


THE  OXIDES  AND  ACIDS  OF  SULPHUR  279 

are  insoluble  in  water,  and  they  are  formed,  of  course,  by  double 
decomposition  from  soluble  salts,  for  example, 

BaCl2  +  Na2C03  =  BaCO3  +  2NaCl 

But  these  salts,  in  general,  dissolve  in  acids  and  consequently  are 
not  precipitated  if  an  acid  is  present.  It  is  for  this  reason  that 
nitric  acid  is  added  to  the  solution  before  making  the  test  for  a 
sulphate  ion. 

309.  Uses  of  Sulphuric  Acid. — Many  uses  to  which  the  acid  is 
put  will  be  described  in  detail  throughout  this  book,  only  a  few  of 
the  more  important  ones  will  be  mentioned  here.     In  the  inor- 
ganic industries  it  is  used  in  large  quantities  in  the  manufacture 
of  other  acids,  salts,  fertilizers,  and  sodium  carbonate,  and  in  the 
refining  of  certain  metals,  such  as  silver  and  copper.     Sulphuric 
acid  is  used  in  large  quantities  in  the  industries  based  on  organic 
chemistry,  such  as  the  manufacture  of  dyes,  nitroglycerine  for 
dynamite,  nitrocellulose   for   smokeless   powders,  trinitrotoluene 
(T.N.T.)  for  shells  and  bombs,  and  in  the  purification  of  crude 
petroleum  in  making  gasoline  and  kerosene. 

310.  Thiosulphates. — Sodium  thiosulphate  is  the  "  hypo  "  of 
the  photographer.     When  sulphur  is  added  to  a  boiling  solution 
of  sodium  sulphite  the  two  substances  slowly  react  and  sodium 
thiosulphate  is  formed : 

Na2SO3  +  S  =  Na2S2O3 

The  reaction  recalls  that  by  which  sodium  sulphite  is  converted 
into  sodium  sulphate  by  oxygen.  The  relationship  between  the 
formulas  of  the  thiosulphates  and  sulphates  led  to  the  name 
selected  for  the  former;  the  replacement  of  one  oxygen  atom  in  a 
sulphate  by  a  sulphur  atom  yields  a  thiosulphate.  The  prefix 
used  in  this  case  is  derived  from  the  Greek  word  for  sulphur. 

Sodium  thiosulphate,  from  which  the  other  salts  are  prepared, 
is  a  by-product  in  the  manufacture  of  sodium  carbonate.  It  is 
consequently  sold  at  a  low  price  and  has  some  important  applica- 
tions. The  product  on  the  market  is  a  hydrate  of  the  formula 
Na2S2O3,  5H2O  and  crystallizes  well  from  water,  in  which  it  is  very 
soluble.  The  salt  was  formerly  called  sodium  hyposulphite,  but 
the  use  of  the  name  was  not  in  accord  with  systematic  chemical 


280  INORGANIC  CHEMISTRY  FOR  COLLEGES 

nomenclature  and  it  was  abandoned  by  chemists  when  the  true 
hyposulphite  was  discovered. 

Free  thiosulphuric  acid  has  never  been  isolated  on  account 
of  the  fact  that  it  decomposes  when  set  free  from  its  salts.  Decom- 
position takes  place  as  indicated  by  the  following  equations: 


Na2S2O3  +  2HC1  =  2NaCl  + 

H2S2O3  =  H2O  +  SO2  +  S 

The  solution  smells  strongly  of  sulphur  dioxide  and  has  a  milky 
appearance  owing  to  the  sulphur  which  separates  as  an  exceedingly 
finely  divided  white  solid. 

One  of  the  chief  uses  of  sodium  thiosulphate  is  as  an  "  anti- 
chlor  "  in  bleaching.  It  is  readily  oxidized  to  sodium  sulphate  by 
free  chlorine  and,  consequently,  removes  the  excess  of  chlorine 
present  after  the  latter  has  been  used  for  bleaching. 

Sodium  thiosulphate  is  used  in  a  number  of  processes  in  analyt- 
ical chemistry  in  which  iodine  is  a  reagent.  The  substances  react 
according  to  the  following  equation: 

2Na2S2O3  +  I2  =  2NaI  +  Na2S4O6 

The  salt  formed  in  the  reaction  is  called  a  tetrathionate  because  it 
contains  four  sulphur  atoms. 

311.  Persulphates.  —  The  salts  of  persulphuric  acid,  H2S2Og, 
are  active  oxidizing  agents  and  have  recently  been  applied  to  some 
interesting  uses.  Persulphuric  acid  is  so  called  because  it  con- 
tains more  oxygen  than  sulphuric  acid  and  yields  hydrogen  peroxide 
when  it  decomposes.  Sodium  persulphate  is  formed  when  a  con- 
centrated solution  of  sodium  hydrogen  sulphate  is  electrolyzed. 
When  the  ions  of  this  salt,  Na+  and  HSC>4~,  are  discharged  by  the 
current,  the  sodium  liberated  reacts  with  the  water  present  and 
forms  hydrogen  and  sodium  hydroxide;  the  HSCU"  ions  liberated 
unite  and  form  persulphuric  acid,  H2S2Og,  which  interacts  by 
double  decomposition  with  some  of  the  sodium  hydrogen  sulphate 
to  form  sodium  persulphate;  and  this  being  difficultly  soluble 
separates  out. 

The  union  of  two  discharged  HSC>4~  ions  to  form  H2S2Og  is 
entirely  analogous  to  the  formation  of  chlorine  molecules,  C12, 


THE  OXIDES  AND  ACIDS  OF  SULPHUR  281 

when  Cl~  ions  are  discharged.  A  solution  of  persulphuric  acid 
can  be  obtained  by  electrolyzing  an  aqueous  solution  of  sulphuric 
acid  containing  about  50  per  cent  of  the  latter.  It  has  been  shown 
that  the  ammonium  salt  can  advantageously  replace  the  sodium 
salt  in  the  preparation  of  persulphates,  and  large  quantities  are  now 
manufactured  in  this  way. 

Persulphuric  acid  is  not  known  in  the  pure  condition;  when  an 
attempt  is  made  to  prepare  it  from  its  salts  it  decomposes  with  the 
formation  of  either  oxygen  or  hydrogen  peroxide,  depending  on  the 
conditions.  This  fact  leads  to  the  view  that  the  acid  is  related  to 
hydrogen  peroxide,  and  that  in  it  some  of  the  oxygen  atoms  are 
joined  as  they  are  in  the  peroxide.  It  will  be  recalled  that  certain 
acids  are  formed  as  the  result  of  the  union  of  oxides  with  water; 
sulphur  trioxide  and  water,  for  example,  give  sulphuric  acid.  Such 
acids  can  be  broken  down  into  oxides  (anhydrides)  and  water.  It 
is  highly  probable  that  hydrogen  peroxide  reacts  with  anhydrides 
in  a  similar  way,  for  acids  are  known,  the  formulas  of  which  bear 
the  same  relation  to  hydrogen  peroxide  that  the  formulas  of  the 
ordinary  acids  bear  to  water.  For  example,  if  sulphur  trioxide  re- 
acts with  hydrogen  peroxide  as  indicated  by  the  following  equation, 

SO3  +  H2O2  =  H2SO5 

we  would  expect  the  resulting  acid  to  yield  hydrogen  peroxide  as 
the  result  of  decomposition.  An  acid  of  the  above  formula  is 
known,  called  Caro's  acid  from  its  discoverer,  which  is  prepared 
by  the  action  of  hydrogen  peroxide  on  strong  sulphuric  acid.  It  is 
a  powerful  oxidizing  agent.  If  2  molecules  of  sulphur  trioxide 
reacted  with  1  of  hydrogen  peroxide  the  resulting  acid  would 
have  the  formula  H2S2Os,  which  is  that  of  persulphuric  acid. 
The  relationship  indicated  is  probably  true  for  the  decomposition 
of  persulphuric  acid  in  water  solutions  yields  either  hydrogen 
peroxide  or  oxygen,  which  is  probably  formed  from  the  peroxide: 

H2S2O8  +  2H2O  =  2H2SO4  +  H2O2 

This  reaction  is  utilized  in  one  of  the  methods  used  industrially  for 
the  preparation  of  hydrogen  peroxide.  Ammonium  persulphate  is 
treated  with  strong  sulphuric  acid;  the  persulphuric  acid  liberated 
breaks  down  spontaneously  into  hydrogen  peroxide,  which  is 


'282  INORGANIC  CHEMISTRY  FOR  COLLEGES 

obtained  from  the  mixture  by  distillation  under  diminished  pres- 
sure. Persulphuric  acid  and  its  salts  are  oxidizing  agents,  a  fact 
which  leads  to  their  use  in  bleaching.  They  decompose  slowly  in 
solution,  but  rapidly  in  the  presence  of  a  substance  which  they 
can  oxidize.  In  this  decomposition  sulphates,  sulphuric  acid, 
and  oxygen  are  formed. 

Persulphates  are  used  in  "  reducing  "  photographic  negatives. 
When  a  negative  has  been  over  developed  too  much  silver  has  been 
formed  and  the  plate  is  too  dense.  As  much  silver  can  be  removed 
from  the  negative  as  desired  by  putting  it  into  a  solution  of  a  per- 
sulphate. The  latter  slowly  oxidizes  the  metal  and  converts  it 
into  silver  sulphate,  which  is  dissolved  by  the  water  present: 

2Ag  +  K2S2O8  =  K2S04  +  Ag2S04 

EXERCISES 

1.  Write  equations  for  the  reactions  which  take  place  between  the  follow- 
ing substances:    (a)  SO2C12+ H2O,     (6)  concentrated  H2SO4  and   Ag,     (c) 
CaSO3  +  H2SO4,    (d)  Na2S2O7  +  H2O,    (e)  concentrated  H2SO4  and  Hg,    (/) 
concentrated  H2SO4  and  As  as  the  result  of  which  As2O3  is  formed. 

2.  Write  the  graphic  formula  of  the  compound  formed  as  the  result  of 
the  addition  of  HC1  to  SO3,  and  an  equation  for  the  reaction  of  the  product 
with  H2O. 

3.  State  two  ways  in  which  you  could  separate  SO2  from  CO2. 

4.  Sulphurous  acid  is  a  much  stronger  acid  than  carbonic  acid.     What 
would  happen  if  sulphuric  acid  were  added  to  a  solution  containing  sodium 
sulphite  and  sodium  carbonate,  the  amount  of  the  acid  being  less  than  that 
required  to  decompose  either  salt? 

5.  What  would  happen  if  dilute  hydrochloric  acid  were  added  to  a  solu- 
tion of    (a)  sodium  sulphide,    (6)  sodium  thiosulphate,  and    (c)  a  mixture 
of  the  two  salts.     Write  equations  for  all  reactions. 

6.  What  volume  of  oxygen  at  0°  and  760  mm.  will  react  with  a  solution 
containing  100  grams  of    (a)  hydrated  sodium  sulphite,    (6)  the  anhydrous 
salt,  and   (c)  sodium  metabisulphite? 

7.  When  an  excess  of  iodine  is  added  to  a  dilute  solution  of  sulphur  dioxide 
the  latter  is  oxidized  quantitatively  to  sulphuric  acid.     Devise  a  volumetric 
method  for  the  quantitative  determination  of  SO2  based  on  this  reaction. 

8.  What  weight  of  sulphur  must  be  burned  to  produce  enough  SO2  to 
saturate  1  liter  of  water  with  the  gi:s  at  room  temperature? 

9.  Starting  with  Na2SO4,10H2O  write  equations  for  reactions    by  which 
the   following    could    be   prepared:    (a)  NaHSO4,     (6)  Na2S2O7,     (c)  Na2S, 
(d)  Na2S03,   (e)  Na2S2O3. 

10.  Write  an  equation  for  the  reaction  between  Na2SOs  and  Na2S2O8. 


THE  OXIDES  AND  ACIDS  OF  SULPHUR  283 

11.  A  mixture  of  sodium  sulphate  and  salt  was  analyzed  with  the  following 
result:   1.000  gram  of  the  mixture  when  dissolved  in  water  and  treated  with 
barium  chloride  gave  1.315  grams  barium  sulphate.     Calculate  the  percent- 
age of  sodium  sulphate  in  the  mixture. 

12.  (a)  How  many  cubic  feet  of  air  are  necessary  to  convert  1  ton  of  pyrite, 
FeS2,  into  SO2?     (6)  What  is  the  relation  between  the  volume  of  the  SO2 
and  N2  if  no  excess  of  air  is  used?     (1  pound  molecular  weight  occupies  359 
cu.  ft.) 


CHAPTER  XXI 
NITROGEN  AND  THE  ATMOSPHERE 

312.  The  air  has  aroused  in  man  the  greatest  interest  since  the 
earliest  days.  It  was  feared  and  venerated,  and  like  other  things 
in  nature  which  affected  man's  existence  for  good  or  evil  it  was 
deified.  In  early  Grecian  days  sailors  prayed  and  burned  incense 
to  the  God  of  the  Storm  before  setting  out  on  a  voyage.  As  men 
began  to  question  nature  they  appreciated  the  significance  of  the 
air  and  it  became  one  of  the  fundamental  concepts  in  ancient 
philosophy.  Empedocles,  a  Greek,  put  forward  in  the  fifth  cen- 
tury before  Christ  the  view  that  everything  was  made  of  four 
elements,  earth,  air,  fire,  and  water.  This  view  persisted  through 
centuries,  but  was  modified  as  time  went  on  and  the  air  was  more 
and  more  studied.  John  Mayow  (1654-1679)  showed  that  when 
certain  metals  were  heated  in  the  air  they  withdrew  something 
from  it;  Priestley  demonstrated  in  1774  that  the  air  contains 
something  which  is  essential  in  combustion,  and  obtained  this 
important  substance  in  pure  condition  by  heating  the  oxide  of 
mercury;  and,  finally,  Lavoisier  explained  the  part  played  in 
combustion  by  Priestley's  gas,  which  he  called  oxygen,  and  demon- 
strated by  convincing  experiments  that  his  view  was  correct.  It 
can  be  said  with  assurance  that  the  science  of  chemistry  was  born 
when  men  first  began  to  study  experimentally  the  nature  of  the  air. 

In  the  early  history  of  the  science  of  physics,  also,  the  study  of 
air  plays  an  important  part.  The  explanations  first  offered  of 
physical  phenomena  were,  in  general,  like  those  put  forward  in 
the  case  of  chemical  changes.  The  forces  of  nature  were  assigned 
human  attributes;  water  could  not  be  pumped  up  into  a  tube 
higher  than  34  feet,  because  nature  "  abhorred  "  a  vacuum.  Tor- 
ricelli  came  to  the  conclusion  that  the  rise  of  the  water  was  due 
to  the  pressure  of  the  air.  If  this  were  true  the  height  of  a  col- 
umn of  a  heavy  liquid  supported  by  air  would  be  less  than  that  of 

284 


NITROGEN  AND  THE  ATMOSPHERE  285 

a  column  of  water.  Torricelli  tried  the  experiment  with  mercury 
in  1643  and,  as  a  result,  invented  the  barometer.  Pascal  observed 
later  that  the  height  of  a  barometer  on  the  top  of  a  mountain  was 
less  than  at  its  base,  a  fact  that  confirmed  the  view  of  Torricelli, 
for  at  the  top  there  is  less  air  to  support  the  mercury.  Guericke 
invented  the  air-pump  in  1650;  Boyle  at  once  made  one  for  him- 
self and  as  the  result  of  his  experiments  with  it  discovered  the  law 
for  which  his  name  is  famous. 

313.  The  atmosphere  has  been  one  of  the  chief  causes  in  shaping 
the  earth's  surface  as  it  is  to-day.     The  carbon  dioxide  and  water- 
vapor  of  the  air  decompose  rocks,  slowly  converting  them  into  the 
materials  of  which  the  soil  is  composed.     For  this  and  other 
reasons  the  influence  of  the  air  and  its  constituents  is  studied  in 
Geology.     Meteorology — the  science  of  climate  and  the  weather — 
is  primarily  concerned  with  air  from  the  physical  point  of  view. 
The  micro-organisms  in  the  air  bring  about  many  chemical  changes 
through  the  agency  of  fermentation;   the  putrefaction  of  animal 
substances,  and  the  decay  of  vegetable  material,   such  as  the 
rotting  of  wood,  are  examples  of  important  changes  that  are  brought 
about  in  this  way.    Pasteur  was  able  to  prove  the  impossibility  of 
spontaneous  generation — the  production  of  life  without  the  presence 
of  living  organisms — by  carrying  out  experiments  in  the  absence 
of  ordinary  air.     When  the  latter  was  freed  from  all  bacteria, 
molds,  etc.,  and  when  the  medium  in  which  the  experiment  was 
carried  out  was  sterilized,  no  living  thing  was  produced. 

NITROGEN 

314.  A  long  time  before  nitrogen  was  isolated,  it  had  been 
observed  that  air  contains  two  substances,  one  of  which  was  inac- 
tive and  did  not  take  part  in  combustion.     Rutherford  in  Scotland 
in  1772  recognized  the  gas  as  a  distinct  substance,  and  at  about  the 
same  time  Scheele  in  Sweden  showed  how  it  could  be  obtained 
by  burning  substances  in  the  air  in  order  to  remove  the  oxygen; 
the  residue  after  removal  of  the  products  of  combustion  was  the 
substance  we  now  call  nitrogen.     It  was  shown  later  that  the 
element  was  present  in  niter  (potassium  nitrate,  KNOs),  and  this 
fact  led  to  the  selection  of  the  English  name  for  the  gas.     Lavoisier 
proved  that  nitrogen  is  an  elementary  substance. 


286  INORGANIC  CHEMISTRY  FOR  COLLEGES 

315.  Occurrence  of  Nitrogen. — About  four-fifths  of  the  air  is 
free  nitrogen.  We  have,  thus,  an  unlimited  supply  of  this  element 
which  plays  such  an  important  part  in  life-processes.  Up  to 
recent  years,  however,  man  was  unable  to  utilize  the  air  as  a 
source  of  nitrogen  compounds  to  be  used  in  fertilizing  the  soil  and 
for  other  important  purposes,  on  account  of  the  great  chemical 
inactivity  of  the  element.  In  certain  natural  processes,  however, 
free  nitrogen  does  take  part;  scientists  have  only  recently  found 
how  to  reproduce  and  use  these  processes.  The  utilization  of  the 
inert  nitrogen  of  the  atmosphere  for  agriculture,  the  basis  of 
civilization  and  wealth,  is  one  of  the  greatest  triumphs  of  modern 
chemistry.  The  processes  already  developed  and  others  which 
will  be  devised,  no  doubt,  will  be  of  incalculable  value  when  the 
supply  of  nitrogen  compounds  available  on  the  earth  is  exhausted 
or  reduced  to  such  an  extent  that  it  does  not  suffice  to  furnish  the 
world's  requirement. 

Inorganic  compounds  which  contain  nitrogen  are  soluble  in 
water,  and,  as  a  consequence,  we  would  not  expect  to  find  them 
accumulated  in  large  quantities  on  the  earth's  surface  except  under 
unusual  circumstances.  In  certain  arid  regions  there  are  supplies 
of  nitrates.  The  chief  commercial  source  of  these  salts  is  the  sodium 
nitrate  obtained  from  Chile  (Chile  saltpeter) .  Guano  was  formerly 
much  used  as  an  ingredient  of  fertilizers  on  account  of  the  fact  that 
it  contains  a  large  percentage  of  nitrogen  compounds.  Guano  is 
obtained  from  certain  tropical  islands.  It  is  the  dried  excrement 
of  sea  gulls,  penguins,  and  other  aquatic  birds  that  breed  in  mil- 
lions along  the  coast  of  South  America.  Ammonia,  NHs,  and 
nitrates  are  found  in  fertile  soils,  and  although  they  are  present 
in  but  small  proportions  they  are  essential  constituents;  in  this 
form  they  do  not  serve,  however,  as  sources  of  nitrogen  compounds 
for  the  industries.  Coal  contains  compounds  which  yield  ammonia 
on  distillation;  the  manufacture  of  coke  and  illuminating  gas  by 
heating  coal  yields  large  amounts  of  ammonia.  Gas-works, 
by-product  coke  ovens,  and  Chile  saltpeter  are  the  chief  sources 
of  the  world's  supply  of  combined  nitrogen  for  industrial  purposes. 

All  living  things  contain  nitrogen  compounds;  the  so-called 
proteins,  which  are  the  chief  constituents  of  flesh  and  are  present 
in  all  vegetable  matter,  contain  approximately  16  per  cent  of  com- 
bined nitrogen. 


NITROGEN  AND  THE  ATMOSPHERE  287 

316.  Preparation  of  Nitrogen. — The  gas  in  a  comparatively 
pure  condition  can  be  readily  obtained  by  burning  a  substance  in 
air,  and  subsequently  removing  the  products  of  combustion.  This 
is  done  in  the  laboratory  by  burning  phosphorus  under  a  bell-jar, 
the  bottom  of  which  dips  under  water.  Phosphorus  is  selected 
because  it  burns  readily  and  the  product  of  combustion,  phos- 
phorus pentoxide,  is  a  solid  which  dissolves  in  water.  The  process 
removes  oxygen  only  and  the  gas  left  contains  water-vapor,  about 
1  per  cent  of  inert  gases,  and  a  trace  of  carbon  dioxide  and  other 
substances.  If  desirable,  all  of  these  except  the  inert  gases  can 
be  removed  by  the  proper  reagents  from  the  nitrogen,  but  to 
exhibit  the  chemical  inertness  of  the  latter  this  is  not  necessary. 

A  convenient  way  of  removing  oxygen  from  the  air  is  to  pass  it 
over  hot  iron  or  copper,  the  latter  being  preferable  in  the  laboratory 
because  it  is  oxidized  rapidly  at  a  comparatively  low  temperature. 
On  account  of  the  expense  of  the  materials  used,  nitrogen  is  not 
obtained  in  these  ways  for  industrial  purposes.  When  the  oxygen 
is  removed  chemically,  coke  or  coal  is  burned  in  the  air,  and  the 
oxides  of  carbon  produced  are  removed. 

Free  nitrogen  can  be  obtained  from  compounds  of  the  element, 
but  since  the  important  commercial  problem  is  to  transfer  nitrogen 
into  these  compounds,  such  processes  are  not  used  in  chemical 
industry.  All  the  oxides  of  nitrogen  are  reduced  when  heated 
with  the  copper  or  other  metals;  an  example  is  illustrated  by  the 
following  equation: 

2NO  +  2Cu  =  2CuO  +  N2 

Ammonia  can  be  oxidized  by  passing  it  over  hot  copper  oxide: 
2NH3  +  3CuO  =  3H2O  +  N2  +  3Cu 

All  organic  compounds  containing  nitrogen  yield  nitrogen  when 
heated  with  copper  oxide.  The  decomposition  is  the  basis  for  the 
quantitative  analysis  of  such  substances.  The  weight  of  the 
material  burned  and  the  volume  of  the  nitrogen  produced  make  it 
possible  to  calculate  the  percentage  of  nitrogen  present  in  the  com- 
pound. 

The  easiest  way  to  make  pure  nitrogen  in  the  laboratory  is  to 
heat  ammonium  nitrite: 

NH4N02  =  N2  +  2H20 


288  INORGANIC  CHEMISTRY  FOR  COLLEGES 

As  the  nitrite  is  a  very  unstable  substance  it  cannot  be  kept.  In 
making  nitrogen  in  this  way  a  strong  solution  of  sodium  nitrite 
and  ammonium  chloride  is  heated;  the  substances  first  interact 
by  double  decomposition  and  form  ammonium  nitrite, 

NH4C1  +  NaNO2  =  NH4N02  +  NaCl 

and  the  latter  then  decomposes  into  nitrogen  and  water. 

An  important  source  of  nitrogen  for  industrial  purposes  is 
liquid  air.  On  evaporation  of  the  liquid,  nitrogen  first  escapes  and 
then  oxygen;  liquid  air  serves,  therefore,  as  a  comparatively  cheap 
source  of  both  these  gases. 

317.  Physical  Properties  of  Nitrogen. — Nitrogen,  N2,  is  a  color- 
less, odorless  gas;    it  is  slightly  lighter  than  air  (sp.  gr.  0.967); 
1  liter  at  0°  and  760  mm.  weighs  1.2506  grams;  it  is  slightly  soluble 
in  water  (2  volumes  in  100  at  6°).     Liquid  nitrogen  boils  at  — 194°, 
and  freezes  at  —214°;    since  its  critical  temperature  is  —146° 
nitrogen   cannot  be   liquefied    by  pressure  at  ordinary  temper- 
atures;  the  liquid  exerts  such  a  high  pressure  that  it  is  kept  in 
open  vessels. 

318.  Chemical  Behavior  of  Nitrogen. — At  ordinary  tempera- 
tures nitrogen  is  inert;    the  only  chemical  reaction  known  into 
which  it  enters  under  these  circumstances  is  one  brought  about 
through  the  influence  of  bacteria  present  in  nodules  on  the  roots 
of  certain  plants,  such  as  peas  and  beans  and  clover.  The  change  of 
nitrogen  to  nitrates  in  this  way  is  of  great  importance  in  agriculture 
and  is  the  scientific  basis  for  the  rotation  of  crops.     Under  the 
influence  of  the  bacteria  some  of  the  nitrogen  of  the  air  is  con- 
verted  into  proteins  and  the  nitrogen  "  fixed  "  in  this  way  is 
changed  by  other  bacteria  present  into  nitric  acid,  which,  penetrat- 
ing the  soil,  serves  as  a  food  for  the  growing  plant.    This  subject 
has  been  much  investigated,  and  the  bacteria  which  affect  the 
free  nitrogen  have  been  isolated.     Seeds  which  have  been  dipped 
into  a  solution  containing  the  bacteria  inoculate  the  soil  when  they 
are   sown,  and   the  growth  of  the  resulting  plant  is  markedly 
increased. 

With  rise  in  temperature  the  activity  of  nitrogen  increases,  and 
at  very  high  temperatures  it  is  one  of  the  most  active  of  all  the 
elements.  When  an  electric  discharge  passes  through  the  air  the 
nitrogen  and  oxygen  in  the  immediate  vicinity  unite  and  form  nitric 


NITROGEN  AND  THE  ATMOSPHERE  289 

oxide,  NO.  As  a  result  of  this  action  nitric  acid,  formed  from  this 
oxide,  oxygen,  and  water-vapor,  is  produced  in  thunder  storms. 
Nitrogen  unites  with  hydrogen  to  form  ammonia,  NHs,  when  the 
two  gases  are  heated  together.  Both  of  these  reactions  have  been 
used  for  the  so-called  fixation  of  atmospheric  nitrogen;  nitric 
acid  and  ammonia,  the  two  most  important  compounds  of  nitrogen, 
are  now  manufactured  on  the  large  scale  from  free  nitrogen. 

At  the  temperature  of  the  electric  furnace  nitrogen  combines 
with  certain  metals  and  non-metals;  some  of  the  products,  which 
are  called  nitrides,  have  interesting  properties  and  may  become 
articles  of  commerce.  Some  nitrides  react  with  water  to  form 
ammonia.  The  reaction  has  been  studied  in  the  hope  of  finding 
a  new  method  of  making  this  industrially  important  substance. 
When  magnesium  is  burned  in  the  air  the  chief  product  is 
magnesium  oxide,  but  at  the  temperature  produced  in  the  burn- 
ing, some  of  the  metal  unites  with  the  nitrogen  in  the  air  and 
magnesium  nitride,  MgsN2,  is  formed.  When  the  nitride  is  boiled 
with  water,  ammonia  is  produced. 

319.  Uses  of  Nitrogen. — It  has  already  been  pointed  out  that 
free  nitrogen  is  used  in  the  preparation  of  nitric  acid  and  ammonia, 
and  that  certain  plants  utilize  the  gas  in  their  growth. 

An  important  use  of  nitrogen  is  in  the  manufacture  of  nitrogen- 
filled  tungsten  lamps.  Up  to  a  short  time  ago  the  bulbs  of  electric 
lamps  contained  no  gas;  the  air  was  pumped  out  in  order  to 
prevent  the  action  of  oxygen  on  the  filament.  When  tungsten 
replaced  carbon  as  the  material  of  which  the  filament  was  made, 
it  was  found  that  a  black  deposit  formed  on  the  walls  of  the  bulb 
after  the  lamp  had  been  used  for  some  time.  The  deposit  proved 
to  be  tungsten,  which  had  distilled  off  from  the  filament  at  the 
high  temperature  to  which  it  was  heated.  It  was  found  that  by 
filling  the  bulb  with  nitrogen  the  blackening  was  largely  prevented. 


THE  ATMOSPHERE 

320.  The  components  of  the  air  that  are  present  in  almost 
constant  proportions  are  oxygen,  nitrogen,  argon,  and  the  other 
rare  gases.  Carbon  dioxide  is  always  present,  but  since  it  is  the 
product  of  respiration,  the  combustion  of  coal  and  wood,  and  the 


290  INORGANIC  CHEMISTRY  FOR  COLLEGES 

decay  of  vegetable  material,  the  quantity  present  in  air  varies 
greatly  when  the  gas  is  introduced  into  the  air  from  these  sources. 
The  amount  of  water-vapor  in  the  air  fluctuates  between  wide 
limits. 

Air  is  analyzed  by  determining  the  relative  volumes  of  the 
gases  present,  and  the  results  are  expressed  as  percentages  or  as 
parts  by  volume.  Air  freed  from  water-vapor,  carbon  dioxide, 
and  accidental  constituents  contains  approximately  21  per  cent 
oxygen,  78  per  cent  nitrogen,  and  slightly  less  than  1  per  cent  of 
argon. 

Samples  of  air  collected  at  different  parts  of  the  earth's  surface 
have  been  examined  and  the  percentages  of  oxygen  found  never 
varied  from  one  another  more  than  0.2  per  cent.  The  actual  amount 
of  oxygen  and  other  constituents  in  a  sample  of  air  varies,  of 
course,  greatly  with  the  pressure  of  the  atmosphere.  The  several 
constituents  of  the  air  will  now  be  considered  in  some  detail. 

321.  The  Oxygen. — The  part  that  oxygen  plays  in  natural 
processes  has  already  been  discussed  (38).  Life  as  we  know  it  on 
the  earth  centers  around  oxygen.  Whenever  we  move  there  is  an 
expenditure  of  energy  and  this  is  supplied  as  the  result  of  oxidation 
in  the  body.  Our  needs  of  this  essential  factor  in  life  are  well 
looked  after.  There  is  an  ample,  supply,  and  the  body  can  func- 
tion when  the  proportion  of  oxygen  in  the  air  is  much  reduced. 
We  can  live  in  an  atmosphere  in  which  a  candle  will  not  burn. 
Oxygen  converts  animal  and  vegetable  refuse  through  the  agency 
of  micro-organisms  into  innocuous  gases. 

The  amount  of  oxygen  in  air  can  be  determined  readily  by 
introducing  into  a  sample  of  known  volume  some  substance  which 
unites  with  oxygen  at  ordinary  temperatures.  The  experiment  is 
usually  carried  out  by  students  in  the  laboratory  with  the  aid  of 
a  eudiometer,  which  is  a  glass  tube  closed  at  one  end  and  marked 
with  lines  etched  into  the  glass  so  that  the  volume  of  the  contained 
gas  can  be  read.  The  open  end  of  the  tube  is  placed  under  water 
and  a  piece  of  phosphorus  supported  on  the  end  of  a  wire  is  inserted 
into  the  air.  The  volume  of  the  air  in  the  tube  is  read,  and  the 
temperature  of  the  outside  air  and  the  height  of  the  barometer 
are  recorded.  After  standing  a  day  or  longer,  until  all  the  oxygen 
has  reacted  with  the  phosphorus,  the  readings  are  made  again. 
The  decrease  in  volume  equals  the  volume  of  the  oxygen  originally 
present. 


NITROGEN  AND  THE  ATMOSPHERE  291 

322.  The  Carbon  Dioxide. — By  far  the  larger  part  of  the  carbon 
dioxide  in  the  air  is  the  product  of  the  decay  of  vegetable  material, 
and,  as  a  consequence,  the  percentage  of  the  gas  in  the  atmosphere 
is  practically  constant,  except  in  large  cities  or  in  the  neighborhood 
of  factories  where  great  amounts  of  coal  are  burned.  Country 
air  contains  3  parts  per  10,000,  or  .03  per  cent;  the  atmosphere  in 
a  large  city  may  have  as  much  as  7  or  8  parts;  and  in  a  crowded 
room  the  content  of  carbon  dioxide  may  rise  as  high  as  50  parts. 

Air  can  be  readily  freed  from  carbon  dioxide  by  passing  it 
through  a  solution  of  sodium  hydroxide.  If  it  is  desired  to  deter- 
mine whether  carbon  dioxide  is  present  in  a  sample  of  gas,  calcium 
hydroxide  is  used  because  in  this  case  the  carbonate  formed  is 
insoluble  and  precipitates;  it  is,  therefore,  visible.  Barium 
hydroxide  can  be  used  in  the  test  instead  of  calcium  hydroxide. 
In  making  an  air  analysis,  a  large  bottle  (8  or  10  liters)  is  filled  by 
blowing  into  it  with  a  bellows,  and  thus  obtaining  a  sample  of  the 
air  in  the  neighborhood.  A  measured  quantity  of  a  solution  of 
barium  hydroxide  of  known  strength  is  next  added,  and  the  bottle 
closed  and  shaken.  After  the  precipitate  has  settled,  some  of  the 
clear  solution  is  drawn  off  in  a  pipette  and  the  amount  of  barium 
hydroxide  in  it  determined  by  neutralization  with  an  acid.  From 
the  result  obtained,  the  amount  of  the  hydroxide  which  did  not 
react  with  the  carbon  dioxide  can  be  calculated;  and  the  difference 
between  this  quantity  and  that  used  is  a  measure  of  the  amount  of 
carbon  dioxide  in  the  air  analyzed. 

The  decay  of  dead  vegetation  furnishes,  as  has  been  said,  the 
bulk  of  the  carbon  dioxide  in  the  air.  This  decomposition  is 
brought  about  by  oxidation  induced  by  vegetable  organisms  which 
thrive  in  the  presence  of  moisture.  As  a  consequence,  when  we 
wish  to  preserve  wooden  buildings  from  decay  we  paint  them. 
The  oil  contained  in  the  paint  protects  the  wood  from  the  action  of 
oxygen,  and  as  moisture  is  not  readily  absorbed,  the  destructive 
organisms  do  not  thrive.  In  time,  however,  the  oil  itself  is  oxidized, 
the  paint  crumbles,  and  the  wood  is  exposed.  Moisture  plays  an 
important  part  in  this  kind  of  decay;  for  this  reason  telegraph 
poles  are  often  set  in  cement  to  keep  the  ends  in  the  earth  away 
from  the  water  present  in  the  soil.  Wooden  buildings  last  much 
longer  in  dry  climates.  Remains  of  temples  built  of  wood  hun- 
dreds of  years  old  have  been  excavated  in  India. 


292  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  burning  of  coal  introduces  large  amounts  of  carbon  dioxide 
into  the  air.  It  has  been  estimated  that  nearly  one  and  one-half 
billion  tons  of  coal  are  used  each  year.  Forest  fires  also  produce 
vast  quantities  of  carbon  dioxide,  but  these  sources  account  for 
less  than  two-tenths  per  cent  of  the  gas  present  in  the  atmosphere. 

Carbon  dioxide  gets  into  the  air  as  the  result  of  respiration  in 
animals.  Expired  air  contains  about  4  per  cent  of  carbon  dioxide, 
and  as  a  man  breathes  about  12  cubic  meters  per  day,  each  indi- 
vidual on  the  earth  is  producing  about  500  liters  or  1000  grams  of 
carbon  dioxide  daily. 

The  constant  accumulation  of  carbon  dioxide  in  the  air  is 
prevented  by  the  fact  already  emphasized,  that  growing  plants 
convert  the  gas  into  the  organic  materials  of  which  they  are  com- 
posed. Carbon  dioxide  is  also  removed  from  the  atmosphere  in 
rain  and  passes  into  rivers  and  finally  into  the  ocean,  which  con- 
tains more  of  the  gas  than  the  atmosphere.  Carbon  dioxide  in 
water  is  converted  into  carbonic  acid  and  slowly  decomposes 
mineral  matter  in  the  soil,  and  is  in  this  way  converted  into  soluble 
salts  which  are  utilized  by  plants  in  their  growth. 

323.  The  Water-vapor. — The  fact  that  water  is  essential  in 
the  natural  processes  that  take  place  on  the  earth  has  been  empha- 
sized (96) .  The  distribution  of  water  is  brought  about  by  the  fact 
that  it  passes  into  the  air  as  a  gas  and,  consequently,  becomes  dis- 
seminated over  the  entire  surface  of  the  globe.  It  returns  to  the 
liquid  state  as  rain  or  dew  and  thus  furnishes  plant  life  with  one  of 
its  necessary  foods. 

Water-vapor  has  been  shown  recently  to  play  a  very  important 
part  in  ventilation.  Various  theories  have  been  put  forward  as  to 
the  cause  of  the  disagreeable  sensations  experienced  in  a  crowded, 
ill-ventilated  room;  lack  of  oxygen,  excess  of  carbon  dioxide,  and 
poisonous  substances  given  off  from  the  lungs  have,  in  turn,  been 
assigned  as  the  cause.  But  none  of  these,  as  has  been  shown,  is 
adequate.  The  changes  in  the  amount  of  oxygen  in  the  air  do  not 
appreciably  affect  us  provided  the  required  minimum  is  present; 
and  it  has  been  shown  that  a  man  can  live  comfortably  when  the 
amount  of  oxygen  is  reduced  far  below  that  obtainable  in  a  room 
under  ordinary  circumstances.  The  amount  of  oxygen  in  the  air 
on  a  mountain  top  is  far  below  that  in  a  crowded  room  but  the 
lack  causes  no  discomfort.  The  carbon  dioxide  in  a  room  filled 


NITROGEN  AND  THE  ATMOSPHERE  293 

with  people  may  increase  to  50  parts  in  10,000,  but  air  containing 
this  amount  of  the  gas  can  be  breathed  without  ill  effects.  The 
so-called  "  poisonous  "  material  from  the  lungs  is  hypothetical, 
and  its  presence  was  assumed  in  searching  for  an  explanation  of  the 
cause  which  produces  the  effects  of  bad  ventilation.  The  body 
gives  off  substances  of  disagreeable  odor  and  these  produce,  no 
doubt,  marked  psychological  effects.  On  the  other  hand,  experi- 
mentation has  shown  that  the  accumulation  of  water-vapor  as 
the  result  of  respiration  from  the  lungs  and  skin  in  a  crowded  room 
has  a  marked  effect  on  comfort.  The  air  space  between  the  indi- 
viduals, crowded  together,  soon  becomes  saturated  with  water- 
vapor,  and  the  temperature  rises.  We  have  thus  the  conditions 
emphasized  which  prevail  in  the  hot,  muggy  days  of  July;  and 
the  same  discomfort  is  evident.  Experiments  have  been  made 
which  prove  the  correctness  of  this  theory;  in  one  of  them  a  num- 
ber of  men  remained  in  a  small  closet,  tightly  closed,  until  they 
suffered  extreme  discomfort  from  "  lack  of  pure  air."  An  elec- 
tric fan  was  then  started;  in  a  short  time  the  moisture  and  other 
constituents  of  the  air  were  evenly  distributed  throughout  the 
closet  and  the  experimenters  lost  the  painful  sensations  they  had 
experienced. 

Although  carbon  dioxide  is  not  the  cause  of  the  effects  pro- 
duced by  bad  ventilation,  the  amount  present  in  the  air  is  usually 
determined  when  a  study  of  the  ventilation  in  a  room  is  being 
made.  This  is  done  because  it  has  been  shown  by  experience  that 
the  amount  of  carbon  dioxide  present  is  a  rough  measure  of  the 
state  of  the  air;  when  it  is  high  the  room  is  uncomfortable,  and  if 
it  is  low  the  fact  is  evidence  that  a  sufficient  quantity  of  fresh  air 
is  being  admitted.  One  of  the  simplest  ways  of  removing  the 
excess  moisture  and  keeping  a  room  at  the  proper  temperature  is 
to  admit  freely  fresh  air,  and  this  method  is  the  one  used  in  proper 
ventilation.  Electric  fans  are  coming  more  into  use  and  are  val- 
uable aids. 

The  facts  stated  above  are  the  basis  for  a  comparatively  new 
science.  The  engineers  who  plan  the  ventilation  of  large  audience 
chambers  take  into  account,  in  estimating  the  amount  of  air  to 
be  admitted  per  minute,  the  number  of  people  the  room  will  accom- 
modate and  the  amount  of  oxygen  consumed,  carbon  dioxide  given 
off,  and  heat  generated  per  individual.  The  humidity  of  the  air 


294  INORGANIC  CHEMISTRY  FOR  COLLEGES 

supplied  is  also  tinder  control.  In  winter  the  amount  of  water- 
vapor  in  the  air  is  small,  because  at  low  temperatures  the  vapor 
pressure  of  water  is  low.  When  such  air  is  heated  it  is  too  dry  for 
comfort;  it  must  be  brought  into  contact  with  water  after  being 
warmed,  so  that  it  can  take  up  what  is  required  to  produce  the 
relative  humidity  desired.  It  would  lead  too  far  to  consider  all 
these  factors  in  detail,  but  the  more  important  principles  upon 
which  ventilation  is  based  are,  no  doubt,  clear. 

324.  Other  Constituents. —  Hydrogen  peroxide  is  formed  in 
minute  quantities  in  the  air.  The  way  in  which  it  is  produced  is  not 
well  understood,  and  there  is  opportunity  for  further  study  of  the 
problem.  It  is  possible  that  in  natural  processes  it  plays  some  part 
which  has  not  been  yet  discovered.  Hydrogen  peroxide  has  been 
found  in  rain-water  and  snow;  it  may  be  produced  as  the  result 
of  electric  disturbances  in  the  atmosphere.  Hydrogen  peroxide 
is  said  to  be  formed  by  the  action  of  sunlight  on  the  surface  of  the 
ocean.  It  is  produced  in  traces  when  zinc,  lead,  or  copper  rust  in 
moist  air.  Hydrogen  peroxide  is  unstable  and  soon  decomposes  in 
the  presence  of  dust  and  bacteria. 

Ozone  is  present  in  the  air  after  thunder-storms.  The 
natural  processes  by  which  it  is  formed  are,  perhaps,  similar  to 
those  which  produce  hydrogen  peroxide.  It  is  said  to  be  a  normal 
constituent  of  air  near  the  sea  and  in  forests,  but  adequate  experi- 
mental evidence  is  lacking  on  this  point.  As  far  as  is  known  it 
plays  no  significant  part  in  nature. 

Hydrogen  occurs  in  the  air,  but  the  amount  present  is  exceed- 
ingly small,  not  more  than  1  part  in  1,500,000.  The  gas  issues 
from  volcanoes,  and  is  formed  as  the  result  of  putrefaction  of 
organic  matter  brought  about  by  certain  bacteria. 

Nitric  acid  is  formed  from  water  and  the  oxides  of  nitrogen  pro- 
duced as  the  result  of  electric  discharges  in  thunder-storms.  The 
acid  dissolves  in  the  rain  and  thus  enters  the  soil  and  is  a  factor  in 
supplying  the  combined  nitrogen  needed  by  plant  life.  The 
ammonia  in  the  air  is  one  of  the  products  of  the  decay  of  organic 
material.  It  also  is  returned  to  the  soil  through  the  agency  of  the 
rain.  It  has  been  estimated  that  about  6  pounds  of  combined 
nitrogen  per  acre  are  obtained  from  the  air  per  year. 

Sulphur  dioxide  gets  into  the  air  chiefly  as  the  result  of  the 
burning  of  coal.  In  large  cities  its  presence  is  of  importance,  for 


NITROGEN  AND  THE  ATMOSPHERE  295 

it  is  converted  into  sulphuric  acid,  which  attacks  materials  of 
construction  made  of  metal  if  they  are  not  well  protected.  The 
effect  is  generally  noticeable  near  large  railroad  stations. 

325.  Dust  is  found  in  all  air  under  normal  conditions.     Its 
presence  in  the  atmosphere  is  the  cause  of  the  formation  of  clouds 
and  fog.     Water- vapor  will  not  change  into  a  liquid  unless  it 
comes  in  contact  with  a  surface  upon  which  it  can  condense. 
When  moist  air  is  cooled  to  the  temperature  at  which  it  is  sat- 
urated with  water-vapor,   condensation  takes  place  on  the  par- 
ticles of  dust  and  a  cloud  is  formed.     If  the  temperature  continues 
to  fall,  the  minute  drops  of  water  in  the  cloud  grow  larger  as  the 
result  of  the  condensation  of  more  water,  and  finally  reach  such  a 
size  that  they  can  no  longer  stay  suspended  and  fall  as  drops  of  rain. 

The  organic  dust  in  the  air  contains  micro-organisms  such  as 
bacteria  and  spores  of  fungi  and  molds.  These  several  organisms 
bring  about  different  kinds  of  fermentation;  one  causes  the  forma- 
tion of  wine  from  grape  juice  as  the  result  of  the  production  of 
alcohol,  and  another  converts  cider  into  vinegar,  which  contains 
acetic  acid.  Putrefactive  bacteria  decompose  animal  material, 
and  poisonous  products  and  gases  with  a  foul  odor  are  produced. 
The  air  also  carries  at  times  pathogenic  bacteria  which  serve  to 
transmit  disease.  Many  of  the  organisms  in  the  air  thrive  best  in 
damp  places  away  from  direct  sunlight.  This  is  one  reason  for 
having  health  resorts  at  high  altitudes  in  a  sunny  climate. 

Air  can  be  readily  freed  from  dust,  both  mineral  and  organic, 
by  being  drawn  through  a  layer  of  raw  cotton.  The  medium  in 
which  bacteria  are  grown  in  experimental  work  is  protected  from 
organisms  of  the  air  by  placing  a  bit  of  cotton  wool  in  the  mouth 
of  the  test-tube  containing  the  material  under  investigation. 

The  presence  of  argon  and  helium  in  the  air  has  been  men- 
tioned. These  gases  are  of  such  interest  and  their  study  has 
led  to  such  important  results,  that  they  deserve  a  somewhat  detailed 
consideration. 

ARGON  AND  RELATED  GASES 

326.  In  describing  argon  and  the  other  gases  like  it  in  the  air  it 
seems  advisable  to  treat  the  subject  historically  and  outline  the 
researches  which  led  to  such  interesting  results.     In  this  way  the 


296  INORGANIC  CHEMISTRY  FOR  COLLEGES 

student  will  become  acquainted  with  the  spirit  which  guides  scien- 
tific research  and  the  methods  used  in  the  attack  of  a  problem. 

Lord  Rayleigh,  a  physicist,  undertook  to  determine  the  density 
of  certain  gases  with  greater  accuracy  than  had  been  attained 
before.  When  a  property  of  -a  substance  is  to  be  determined 
accurately  it  is  necessary,  of  course,  to  have  it  in  the  purest  con- 
dition possible.  In  order  to  avoid  the  possibility  of  the  presence 
of  unknown  accidental  impurities  which  may  not  be  removed  by 
the  processes  of  purification  used,  it  is  advisable  to  study  samples 
of  the  substance  obtained  from  widely  different  sources  in  which 
the  accidental  impurities  may  be  different. 

Lord  Rayleigh  took  the  precaution  to  weigh  with  great  accu- 
racy nitrogen  obtained  from  the  air  and  that  formed  as  the  result 
of  the  decomposition  of  certain  compounds  containing  the  element. 
The  atmospheric  nitrogen  when  freed  from  all  known  substances 
was  weighed  in  a  bulb  which  contained  0.14332  gram  of  the  gas 
measured  under  standard  conditions.  When  the  same  bulb  was 
filled  with  nitrogen  obtained  from  ammonium  nitrite,  the  gas  under 
the  same  conditions  weighed  0.14256  gram.  This  difference  was 
very  small,  but  considerably  more  than  the  experimental  error, 
and  it  was  always  found  as  the  result  of  repeated  experiments. 
In  the  case  of  the  other  gases  studied  similar  results  were  not 
obtained;  the  weight  of  each  gas  was  constant  whatever  its 
source. 

327.  The  next  step  was  to  explain  the  facts,  and  as  the  ques- 
tion was  a  chemical  one  Professor  Ramsay,  a  chemist,  was  asked 
to  co-operate  in  the  investigation.  Various  hypotheses  were  put 
forward:  One  or  the  other  or  both  samples  of  nitrogen  might  be 
impure;  the  atmospheric  nitrogen  might  not  have  been  freed 
completely  from  known  substances  or  it  might  contain  a  new  sub- 
stance, heavier  than  pure  nitrogen,  which  was  not  removed  in  the 
processes  used;  or  the  nitrogen  from  the  ammonium  nitrite  used 
in  the  preparation  of  the  gas  might  contain  a  trace  of  hydrogen. 
These  suggestions  were  tested  one  after  another,  and  it  was  found 
that  all  known  substances  had  been  eliminated;  the  logical  con- 
clusion was  that  the  air  contained  a  very  small  amount  of  an 
unknown  gas  heavier  than  nitrogen. 

The  next  problem  was  to  isolate  the  gas.  To  do  this  nitrogen 
containing  the  unknown  substance  was  repeatedly  passed  over 


NITROGEN  AND  THE  ATMOSPHERE  297 

heated  magnesium  which  united  with  the  nitrogen  and  therefore 
removed  it.  The  residual  gas  had  a  volume  about  1  per  cent  of 
that  of  the  nitrogen  used.  Nitrogen  was  also  removed  by  a 
method  that  had  been  used  by  Cavendish  in  1785.  The  gas  was 
sparked  in  the  presence  of  oxygen;  under  these  circumstances  the 
nitrogen  was  converted  into  nitric  acid,  which  was  removed  by 
dissolving  it  in  a  solution  of  sodium  hydroxide.  The  excess  of 
oxygen  was  taken  up  by  passing  the  gas  over  hot  copper.  It  is 
worthy  of  note  here  that  when  Cavendish  made  his  experiment 
he  found  that  he  always  obtained  a  residue  which  could  not  be 
removed  by  continued  sparking,  and  he  was  unable  to  explain 
his  results.  He  states  that  the  volume  of  the  residual  gas  was 
T|-O  of  the  nitrogen  used — a  figure  which  agrees  well  with  the 
results  of  the  experiments  carried  out  over  a  century  later. 

The  properties  of  the  gas  obtained  in  this  way  were  care- 
fully studied.  It  was  found  to  be  twenty  times  as  heavy  as 
hydrogen,  whereas  nitrogen  is  fourteen  times  as  heavy.  This 
fact  accounted  for  the  difference  in  weight  of  the  nitrogen  sam- 
ples observed  by  Rayleigh. 

328.  The  gas  was  evidently  a  very  stable  substance,  since  it  had 
resisted  very  active  chemical  reagents  and  high  temperatures  in 
the  course  of  its  isolation.     It  was  thought  it  might  have  been 
formed  from  the  constituents  of  the  air  as  the  result  of  these  vig- 
orous processes;  it  might,  for  example,  be  a  polymer  of  nitrogen  or 
some  unknown  compound.     To  test  this  view  nitrogen  was  freed 
from  known  gases  at  a  low  temperature  and  allowed  to  pass 
slowly  through  unglazed  clay  pipes  surrounded  by  a  vacuum.     It 
will  be  recalled  that  heavy  gases  diffuse  more  slowly  than  light 
ones  (50) .     The  gas  was  passed  repeatedly  through  the  clay  pipes 
and  each  time  a  part  of  it  was  lost  by  diffusion  through  the  walls 
of  the  pipes  into  the  vacuum.     Samples  of  the  gas  were  weighed 
from  time  to  time  and  it  was  found  that  they  increased  in  weight; 
the  lighter  nitrogen  diffused  more  rapidly  than  the  heavier  gas, 
which,  accordingly,  became  more  concentrated.     It  was  evident 
from  these  experiments  that  the  new  gas  was  present  in  the  air 
and  was  not  formed  as  the  result  of  any  chemical  reaction  taking 
place  in  the  process  of  its  isolation. 

329.  The  chemical  properties  of  the  gas  were  then  studied  and 
it  was  found  that  it  entered  into  chemical  reaction  with  nothing. 


298  INORGANIC  CHEMISTRY  FOR  COLLEGES 

It  was  subjected  to  the  most  active  reagents;  metals,  sodium 
hydroxide,  sodium  peroxide,  sodium  and  calcium  persulphides, 
all  red  hot,  did  not  affect  the  gas.  Nascent  chlorine  and  fluorine, 
the  most  active  of  all  the  elements,  failed  to  affect  it.  It  with- 
stood the  silent  electric  discharge  and  the  electric  arc.  There  was 
no  known  element  that  would  not  enter  into  chemical  reaction 
under  many  of  the  conditions  to  which  the  new  gas  was  subjected. 
This  unique  behavior  led  to  the  selection  of  the  name  argon  for  the 
gas,  the  word  being  derived  from  the  Greek  word  signifying 
inactive  or  idle.  The  atomic  weight  of  argon  was  found  to  be  40 
and  it  was  shown  to  have  one  atom  only  in  the  molecule. 

330.  No  use  was  found  for  argon  for  a  number  of  years,  but  its 
inertness  and  the  fact  that  it  is  a  monatomic  gas  finally  led  to 
an  important  application.  The  part  that  nitrogen  plays  in  the 
gas-filled  tungsten  lamp  has  already  been  explained  (319).  While 
the  presence  of  nitrogen  prevents  the  blackening  of  the  globe  by 
the  material  volatilized  from  the  filament,  it  was  found  that  the 
lamp  grew  hot  and,  as  a  result,  an  undue  proportion  of  the  elec- 
trical energy  was  converted  into  heat.  The  amount  of  heat  lost 
from  the  lamp  depended,  evidently,  on  the  rate  at  which  the  heat 
was  conducted  away  from  the  filament.  Heavy  molecules  in  the 
gaseous  condition  move  more  slowly  than  light  ones  (50);  as  a 
consequence,  the  rates  at  which  gases  conduct  heat  vary  with  the 
densities  of  the  gases.  The  replacement  of  nitrogen,  which  has 
the  molecular  weight  28,  by  argon  (40)  reduced  the  rate  at  which 
heat  energy  was  lost  from  the  lamp.  Another  reason  for  the  re- 
placement of  nitrogen  by  argon  is  based  on  the  fact  that  the 
heat  required  to  raise  to  the  same  extent  the  temperatures  of 
equal  volumes  of  the  two  gases,  is  much  less  in  the  case  of  argon. 
The  difference  is  traceable  to  the  monatomic  structure  of  the 
inert  gas.  Argon  was  substituted  for  nitrogen  and  a  more 
efficient  lamp  was  the  result — that  is,  a  greater  proportion 
of  electrical  energy  was  transformed  into  light  because  less 
heat  was  lost.  The  argon  required  was  obtained  from  liquid 
air. 

We  shall  have  more  to  say  of  these  lamps  later  when  tungsten  is 
considered.  They  represent  the  combination  of  the  results  of 
many  researches  in  widely  different  fields — researches  inspired  by 
different  motives.  Rayleigh's  desire  to  weigh  nitrogen  accurately, 


NITROGEN  AND  THE  ATMOSPHERE  299 

Scheele's  study  of  the  mineral  now  known  as  scheelite,  which 
resulted  hi  the  discovery  of  tungsten,  and  the  work  of  the  physi- 
cists who  devised  the  way  to  liquefy  air,  produced  results  which 
were  available  to  the  inventors  and  engineers  who  produced  the 
most  efficient  means  of  artificial  illumination.  The  researches 
mentioned  were  in  what  is  sometimes  inappropriately  called  pure 
science;  they  were  carried  out  with  the  aim  of  finding  out  the 
facts  and  without  thought  of  their  application.  The  scientists 
who  developed  the  lamp  had  a  definite  object  in  view;  their 
researches  were  in  what  is  called  applied  science.  They  utilized 
the  results  obtained  by  many  other  workers,  and  selecting  here 
and  there,  attained  the  desired  result.  Incidentally  important 
new  facts  and  principles  were  discovered  that  may  prove  valuable 
later. 

The  methods  of  research  are  the  same  in  all  fields;  in  applied 
science,  however,  the  aim  is  very  definite — some  particular  thing 
is  to  be  done  and  the  problem  is  oftentimes  for  this  reason  a  more 
difficult  one  to  solve  than  a  problem  in  pure  science.  It  is  impor- 
tant to  note,  however,  that  the  principles  and  facts  of  pure  science 
are  the  building  blocks  of  applied  science.  In  order  to  do  a  thing 
better  than  it  has  been  done  in  the  past,  or  to  do  a  new  thing,  we 
must  be  able  to  call  to  our  aid  knowledge  that  has  not  been  utilized 
for  this  purpose  in  the  past.  And  it  often  happens  that  things 
considered  highly  theoretical  or  impractical  have  become  of  vital 
importance  in  the  production  of  things  of  practical  value.  The 
more  one  knows  of  facts  and  principles  which  have  not  been  applied, 
the  greater  the  store  to  be  drawn  on  when  new  problems  arise. 

331.  Helium. — In  seeking  a  source  of  argon  other  than  the 
air,  Ramsay  was  led  to  examine  certain  minerals  which  were 
reported  to  yield  free  nitrogen  when  heated.  Among  these  was 
the  mineral  clevite.  The  gas  given  off  was  examined  with  a  spec- 
troscope and  found  to  be  a  mixture  in  which  the  element  helium 
was  present.  Lockyer  in  1863,  in  examining  the  spectrum  of  the 
chromosphere  of  the  sun  during  an  eclipse,  noted  an  orange  line 
which  had  not  been  observed  in  the  spectra  of  anything  of  terres- 
trial origin.  He  drew  the  conclusion  that  the  sun  contains  *an 
element  unknown  on  the  earth,  and  he  called  it  helium,  He, 
deriving  the  name  from  the  Greek  word  meaning  the  sun. 

Ramsay  studied  the  gas  in  the  way  used  with  argon  and  found 


300  INORGANIC  CHEMISTRY  FOR  COLLEGES 

that  it  resembled  it  in  being  inactive  and  monatomic.  It  was, 
however,  a  very  light  gas,  being  only  twice  as  heavy  as  hydrogen 
and  one-tenth  as  heavy  as  argon;  its  atomic  weight  was  found  to 
be  four.  The  world  was  searched  over  by  many  investigators  for 
other  sources  of  helium  and  it  was  discovered  occluded  in  a  number 
of  minerals,  dissolved  in  the  water  of  certain  mineral  springs,  and 
as  a  constituent  of  natural  gas  from  certain  sources.  It  was  found 
later  to  be  present  in  the  air  to  the  extent  of  1.4  parts  per  million. 
Helium  was  liquefied  by  Onnes  by  subjecting  it  to  the  temperature 
obtained  by  boiling  liquid  hydrogen  in  a  vacuum.  It  has  a  lower 
boiling-point  than  that  of  any  other  known  substance  (—268.5° 
or  4.5°  Abs.).  When  helium  was  boiled  in  a  vacuum  a  part 
solidified  at  a  temperature  which  was  estimated  to  be  2°  absolute. 
The  most  interesting  fact  about  helium  is  that  it  is  found  in  the 
gaseous  emanation  given  off  by  radium,  and  is  a  product  of  the 
spontaneous  decomposition  of  the  element.  This  far-reaching 
fact  will  be  discussed  in  some  detail  later,  since  it  has  led  to  a  new 
conception  of  the  constitution  of  matter  and  markedly  enlarged 
our  knowledge  of  the  genesis  of  the  earth  and  the  materials  of 
which  it  is  composed, 

332.  We  have  in  helium  another  striking  example  of  how  a  sub- 
stance of  theoretical  significance  suddenly  becomes  of  great  prac- 
tical interest.  Again,  it  is  the  combination  of  certain  properties 
that  make  it  valuable.  In  the  case  of  helium  it  is  its  extreme 
lightness  and  chemical  inactivity.  It  was  suggested  during  the 
recent  war  that  helium  could  be  substituted  for  hydrogen  in 
dirigible  and  observation  balloons.  The  gas  has  almost  the  same 
lifting  power  as  hydrogen,  and  being  non-inflammable  could  not 
be  ignited  by  tracer  bullets  which  were  invented  during  the  war 
to  carry  a  flame  to  a  balloon  and  explode  the  hydrogen  which  it 
contained.  Natural  gas  obtained  in  certain  parts  of  Texas  con- 
tains 0.85  per  cent  of  helium.  The  advisability  of  using  helium 
for  war  purposes  was  so  great  that  the  problem  of  extracting  it 
from  a  source  in  which  it  was  present  in  such  small  quantity  was 
undertaken.  The  process  adopted  involved  liquefaction  of  the 
nakiral  gas  and  isolation  of  the  helium  from  the  residue  which  did 
not  liquefy.  The  pure  gas  compressed  in  steel  cylinders  was 
ready  for  transportation  when  hostilities  ceased.  There  is  no 
doubt  that  this  use  for  helium  will  be  developed,  since  great  dan- 


NITROGEN  AND  THE  'ATMOSPHERE  301 

ger  is  avoided  by  replacing  hydrogen  in  balloons  by  a  non-inflam- 
mable gas. 

333.  Neon,  Krypton,  and  Xenon. — Ramsay  continued  his 
search  for  other  gases  and  distilled  a  large  amount  of  liquid  argon 
obtained  from  the  air.  He  found  that  he  had  a  mixture,  and  in 
addition  to  pure  argon  small  quantities  of  helium  and  of  three 
new  gases  were  obtained.  These  resembled  argon  in  being  chem- 
ically inert  and  monatomic.  Greek  words  were  again  used  in 
naming  the  gases.  Neon  (new)  has  the  atomic  weight  20.2, 
krypton  (hidden)  82.9,  and  xenon  (stranger)  130.2.  Krypton 
is  of  interest  because  the  lines  of  its  spectrum  have  been  observed 
in  the  aurora  borealis  or  northern  lights.  No  adequate  explana- 
tion of  this  striking  natural  phenomenon  has  yet  been  given  and 
there  is  no  reason  known  why  krypton  should  be  present  in  quan- 
tities sufficient  to  make  its  spectrum  visible. 

The  number  of  volumes  of  air  which  contain  one  volume  of  the 
inert  gases  are  as  follows:  Argon,  106.8;  neon,  80,800;  helium, 
245,000;  krypton,  20,000,000;  xenon,  170,000,000. 

EXERCISES 

1.  (a)  Calculate  the  weight  of  1  liter  of  He.     (6)  Calculate  the  relative 
lifting  power  of  helium  and  hydrogen  when  used  in  balloons. 

2.  (a)  How  could  you  determine  the  amount  of  air  exhaled  per  day  by 
a  man  at  rest?     (6)  Make  the  experiment  in  the  laboratory  and  calculate 
the  amount  of  carbon  dioxide  produced  in  twenty-four  hours  assuming  that 
exhaled  air  contains  4  per  cent  CO2. 

3.  A  sample  of  air  was  analyzed  for  carbon  dioxide  with  the  following 
results:  10  liters  of  air  were  shaken  in  a  bottle  with  100  c.c.  of  0.1N  Ba(OH)2. 
After  the  precipitated  BaCO3  had  settled  50  c.c.  of  the  clear  liquid  was  drawn 
off  and  neutralized  with  a  0.1N  solution  of  an  acid,  29.10  c.c.  of  the  solution 
being  required.     Calculate  the    (a)  weight  of  CO2  in  the  sample  of  air,    (6) 
percentage  by  volume  of  CO2  in  the  air,  and    (c)  the  parts  CO2  in  1000  by 
volume. 

4.  The  solubility  of  O2  in  water  at  0°  and  76  cm.  is  approximately  4  vol- 
umes in  100  volumes,     (a)  Assuming  air  to  be  made  up  of  f  N2  and  1  O2, 
what  is  the  pressure  of  the  N2  and  O2  in  air?     (&)  What  volume  of  the  two 
gases  would  be  dissolved  in  100  c.c.  of  water  if  air  were  bubbled  through 
the  latter  until  it  was  saturated?     (See  Henry's  law.)     (c)  How  could  you 
determine  by  an  experiment  the  relation  between  the  N2  and  O2  dissolved 
by  water  when  in  contact  with  air?     (d)  If  air  were  a  compound  of  nitrogen 
and  oxygen  what  relation  would  exist  between  the  nitrogen  and  oxygen  in 
the  air  and  in  the  part  soluble  in  water? 

5.  State  two  reasons  for  the  belief  that  air  is  a  mixture  and  not  a  com- 
pound. 


CHAPTER  XXII 
AMMONIA  AND  ITS  DERIVATIVES 

334.  The  chemistry  of  ammonia  is  of  the  greatest  importance 
because  of  its  significance  in  the  growth  and  decay  of  living  things. 
The  farmer  fertilizes  his  fields  with  ammonium  salts  and  later 
when  he  eats  the  grain  or  vegetables  he  has  grown  he  obtains  the 
nitrogenous  material  which  through  digestion  and  assimilation 
becomes  a  part  of  his  flesh.  Since  the  supply  of  ammonia  fur- 
nished by  nature  is  limited,  the  chemist  has  devised  ways  of  con- 
verting the  nitrogen  of  the  air  into  this  important  product,  and  has 
triumphed  over  what  appeared  at  first  insurmountable  difficulties 
in  his  endeavor  to  make  out  of  the  inexhaustible  supply  of  air 
and  water  a  substance  that  could  be  used  as  a  plant  food  when  the 
naturally  occurring  nitrogen  fertilizers  are  exhausted. 

Since  ammonia  is  produced  as  the  result  of  the  decay  of  refuse 
organic  matter,  its  presence  in  undue  amounts  in  natural  water  is 
evidence  of  contamination.  As  a  consequence,  in  examining  a 
water  supply  to  be  used  for  drinking  purposes,  the  amount  of 
ammonia  present  is  always  determined;  the  results  obtained 
guide  the  chemist  in  his  decision  as  to  its  potability  and  may  lead 
to  the  discovery  of  sources  of  contamination  which  can  be  removed. 

The  physics  of  ammonia  serves  mankind  in  another  important 
way.  The  gas  can  be  readily  liquefied  through  pressure  alone 
and  the  resulting  liquid,  boiling  at  a  very  low  temperature,  absorbs 
a  large  amount  of  heat  when  it  passes  into  gas — its  latent  heat  of 
vaporization  is  very  high.  These  facts  have  been  utilized  in 
making  machines  for  the  production  of  low  temperatures  (180). 
By  the  use  of  ammonia  we  can  make  ice  anywhere  in  unlimited 
quantity,  or  keep  a  warehouse  cold  for  the  storage  of  perishable 
food  materials.  This  application  of  ammonia  has  had  a  marked 
effect  on  the  health  of  people  and  has  revolutionized  the  economics 
of  food  production;  refrigerator  cars  and  cold-storage  warehouses 

302 


AMMONIA  AND  ITS  DERIVATIVES  303 

have  become  a  necessity  in  modern  civilization  with  the  growth  of 
large  cities. 

335.  History    and    Occurrence    of    Ammonia. — Ammonia    is 
formed  when  animal  matter  decays,  and  as  it  has  a  characteristic 
odor  it  has  been  known  from  the  earliest  times.     Ammonium 
chloride — the  compound  formed  from  ammonia  and  hydrochloric 
acid — was  known  to  the  Egyptians.     Its  ancient  name,  sal  ammo- 
niac, from  which  the  word  ammonia  is  derived,  is  said  to  have  been 
given  to  the  salt  in  honor  of  the  Egyptian  sun  god  Ammon.     In 
the  Middle  Ages  ammonia  gas  was  made  by  distilling  the  horns 
of  harts,  and  was  supposed  to  be  very  valuable  as  a  medicinal 
agent.     It  was  called  spirits  of  hartshorn,  a  name  by  which  it  was 
known  until  very  recently  when  used  in  medicine.     The  alchemists 
obtained  ammonium  chloride  by  evaporating  to  dryness  and 
heating  to  a  high  temperature  a  mixture  of  urine  and  salt. 

Priestley,  in  1774,  first  isolated  ammonia  as  a  gas  as  the  result 
of  heating  ammonium  chloride  with  lime.  He  collected  the  gas 
over  mercury.  Many  of  Priestley's  most  important  discoveries 
were  due  to  the  fact  that  he  used  mercury  in  his  pneumatic  trough 
instead  of  water.  He  isolated  in  this  way  for  the  first  time,  in 
addition  to  ammonia,  sulphur  dioxide  and  hydrogen  chloride, 
gases  which  could  not  have  been  recognized  if  water  had  been  used. 
Priestley's  most  important  discovery — that  of  oxygen — was  also 
made  in  his  pneumatic  trough  and  gas  tube  filled  with  mercury. 
It  has  often  happened  that  important  discoveries  have  sprung 
from  the  use  of  new  methods  of  carrying  out  chemical  operations  or 
reactions.  One  can  never  foretell  what  may  result  from  such  a 
new  method;  and  the  discoverer  has  added  a  new  tool  to  the  science. 

336.  Preparation   of   Ammonia. — The   commercial   source   of 
ammonia  is  ammonium  salts,  and  the  gas  is  prepared  from  these 
by  the  action  of  a  base;  ammonium  chloride  and  calcium  hydrox- 
ide are  commonly  used  in  the  laboratory.     As  the  result  of  the 
reaction,  which  is    one    of    double    decomposition,   ammonium 
hydroxide  is  first  formed  and  then  decomposes  into  ammonia  and 
water.     The  equations  for  the  reactions  are  as  follows: 

2NH4C1  +  Ca(OH)2  =  2NH4OH  +  CaCl2 
NH4OH  =  NH3  +  H20 


304  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  reaction  is  brought  about  with  the  solid  substances  containing 
only  a  trace  of  water,  on  account  of  the  fact  that  ammonia  is  very 
soluble  in  water.  The  gas  generated  is  collected  in  the  laboratory 
by  downward  displacement  of  air.  The  collecting  vessel  is  held 
with  the  mouth  downward  because  ammonia  is  only  slightly 
more  than  half  as  heavy  as  air  and,  consequently,  rises. 

When  calcium  hydroxide  reacts  with  ammonium  chloride  in 
aqueous  solution  the  double  decomposition  represented  by  the 
equation  given  above  is  brought  about  as  the  result  of  the  fact 
that  ammonium  hydroxide  breaks  down  spontaneously  to  form 
ammonia.  In  general,  ammonia  is  formed  when  a  substance 
that  furnishes  an  ammonium  ion  is  treated  in  solution  with  a 
substance  that  yields  a  hydroxyl  ion. 

The  chief  sources  of  ammonia  are  the  products  obtained  as  the 
result  of  the  distillation  of  coal  to  make  illuminating  gas  or 
coke  for  metallurgical  uses  (226).  The  gases  formed  in  the  dis- 
tillation are  scrubbed  with  water.  The  solution  which  results  is 
heated  with  lime  to  decompose  the  ammonium  salts  formed  as  the 
result  of  the  neutralization  of  a  part  of  the  ammonia  by  the  carbon 
dioxide  and  other  acidic  material  produced.  The  gas  which  escapes 
is  absorbed  in  sulphuric  acid  and  the  resulting  solution  on  evapora- 
tion yields  ammonium  sulphate.  Hydrochloric  acid  is  at  times 
used  instead  of  sulphuric  acid  and  ammonium  chloride  is  obtained. 

Coal  contains  only  about  1  per  cent  of  nitrogen,  and  the 
yield  of  ammonia  is,  consequently,  small,  but  such  large  quan- 
tities of  coal  are  distilled  that  the  supply  of  ammonia  has  been 
adequate.  When  coke  is  made  in  bee-hive  ovens  (224),  all  the 
volatile  materials  are  lost.  As  the  demand  for  ammonia  increased, 
more  and  more  of  these  ovens  were  replaced  by  by-product  ovens 
in  which  these  valuable  materials  are  condensed  and  saved. 

337.  Synthetic  Ammonia. — It  has  been  known  for  many  years 
that  ammonia  is  formed  when  electric  sparks  are  passed  through  a 
mixture  of  nitrogen  and  hydrogen.  The  behavior  of  nitrogen  with 
hydrogen  under  these  circumstances  is  quite  different  from  that  of 
oxygen.  A  single  spark  is  all  that  is  required  to  cause  the  union 
of  the  entire  amount  of  oxygen  provided  enough  hydrogen  is 
present;  the  reaction  propagates  itself  through  the  entire  mass. 
We  could  synthesize  water  from  oxygen  and  hydrogen  in  any 
desired  quantity  if  it  were  necessary.  With  nitrogen,  however, 


AMMONIA  AND  ITS  DERIVATIVES  305 

only  the  gases  in  the  immediate  vicinity  of  the  flame  unite  and 
the  mixture  has  to  be  sparked  for  a  long  time  to  get  appreciable 
quantities  of  ammonia. 

338.  The  reaction  between  nitrogen  and  hydrogen  has  been 
studied  exhaustively  with  the  hope  of  working  out  a  method  for 
the  synthesis  of  ammonia  which  would  be  of  commercial  value. 
Ammonia  itself  is  a  necessity,  but  the  fact  that  it  can  be  readily 
oxidized  to  nitric  acid  makes  its  commercial  synthesis  of  added 
importance.     The  supply  of  nitric  acid  is  largely  obtained  from 
Chile  saltpeter,  and  as  this  is  not  inexhaustible  chemists  have  been 
studying  in  recent  years  the  "  fixation  of  atmospheric  nitrogen." 
Since  nitric  acid  is  involved  in  making  the  explosives  used  in 
warfare,  and  as  the  sodium  salt  is  an  important  ingredient  of  fer- 
tilizers, the  governments  of  several  nations  have  recently  sub- 
sidized the  study  of  this  important  problem.     During  the  recent 
war  the  blockade  of  Germany  prevented  the  importation  of  salt- 
peter from  Chile  and,  as  a  consequence,  large  amounts  of  ammonia 
were  synthesized  and  converted  into  nitric  acid.     The  Germans,  in 
all  probability,  had  foreseen  the  possibility  of  a  blockade  in  case  of 
war  and  had  bent  their  energies  to  devise  a  commercial  synthetic 
method  for  the  preparation  of  ammonia.     Nitric  acid,  we  shall 
soon  see,  can  be  made  directly  from  the  air  and  water,  but  the 
method  used  requires  large  amounts  of  electrical  energy  and  can 
be  employed  only  where  the  energy  is  cheap  as  the  result  of  the 
availability  of  water-power.     The  production  of  nitric  acid  from 
ammonia  and  the  synthesis  of  the  latter  itself  do  not  require 
electrical  energy,  and,  therefore,  this  method  is  preferred  where  elec- 
tricity must  be  made  by  burning  coal.     The  pressing  need  of 
Germany  was  finally  solved  by  Haber,  and  the  process  of  making 
ammonia  from  nitrogen  and  hydrogen  is  known  by  his  name. 

339.  We  shall  first  examine  carefully  the  reaction  by  which 
ammonia  is  formed  from  nitrogen  and  hydrogen,  and  then  see 
how  it  was  put  on  a  manufacturing  basis.     The  study  of  this  case 
will  bring  out  clearly  how  the  chemist  is  able  to  modify  the  con- 
ditions under  which  a  reaction  takes  place  and  thus  make  it  pos- 
sible to  obtain  the  desired  result. 

In  the  application  of  a  chemical  reaction  to  the  preparation  of  a 
substance  it  is  necessary  to  know  first  whether  the  reaction  is  a 
reversible  one.  If  this  is  the  case  the  equilibrium  attained  in  the 


306  INORGANIC  CHEMISTRY  FOR  COLLEGES 

reaction  under  different  conditions  must  be  determined  in  order 
to  get  the  largest  possible  amount  of  the  desired  product.  The 
equilibrium  in  a  reversible  reaction  is  affected  by  the  temperature 
and,  in  many  cases,  by  pressure,  and  the  rate  at  which  it  is  reached 
is  modified  by  catalyzers.  If  a  chemical  reaction  involving  gases 
is  to  be  intelligently  utilized  it  must  be  studied  from  the  above 
point  of  view  and  a  large  amount  of  data  accumulated,  by  a  con- 
sideration of  which  the  most  efficient  conditions  can  be  arrived 
at.  The  reaction  by  which  ammonia  is  synthesized  is  as  follows: 

N2  +  3H2  *±  2NH3  +  2  X  12,200  cals. 

The  amount  of  heat  developed  must  be  determined  because  the 
reaction  is  a  reversible  one  and  the  equilibrium  attained  is  deter- 
mined by  the  temperature.  In  planning  the  apparatus  in  which 
the  reaction  is  to  be  carried  out,  provision  must  be  made  to  keep 
the  materials  at  the  proper  temperature;  this  is  done  by  with- 
drawing a  part  of  the  heat  formed  or  by  adding  more  heat  as  the 
case  requires. 

The  equilibria  with  change  in  temperature  must  be  deter- 
mined, and  since  the  reaction  is  exothermic  in  this  case,  rise  in 
temperature  is  accompanied  by  shifting  of  the  equilibrium  to  pro- 
duce a  greater  percentage  of  nitrogen  and  hydrogen  (van't  HofFs 
law,  267).  The  reaction  must  then  be  run  at  as  low  a  temperature 
as  is  consistent  with  the  production  of  a  practical  amount  of 
ammonia.  The  lowering  of  the  temperature  has  the  undesired  effect 
of  slowing  up  the  rate  of  reaction.  This  result  was  overcome  by 
finding  a  suitable  catalyzer  to  increase  the  rate.  It  is  seen  from 
the  equation  that  the  gases  react  in  the  relation  of  1  molecule 
of  nitrogen  to  3  of  hydrogen  and  form  2  molecules  of  ammonia. 
We  have  seen  (77)  that  this  indicates  that  1  volume  of  nitrogen 
unites  with  3  volumes  of  hydrogen  and  forms  2  volumes  of 
ammonia;  there  is  a  contraction  from  4  volumes  to  2.  When 
there  is  a  change  in  volume  as  the  result  of  a  chemical  reaction  the 
pressure  under  which  it  takes  place  has  an  effect  on  the  equilibrium 
(law  of  Le  Chatelier,  269).  The  result  is  what  we  would  expect — 
increased  pressure  favors  a  reaction  which  leads  to  decreased 
volume.  It  is  evident  that  a  larger  percentage  of  nitrogen  and 
hydrogen  is  converted  into  ammonia  as  the  pressure  under  which 
the  reaction  takes  place  is  increased;  the  amount  of  pressure  to 


AMMONIA  AND  ITS  DERIVATIVES  307 

be  used  is  determined  by  the  engineering  difficulties  encountered 
in  the  construction  of  the  apparatus. 

In  order  to  determine  the  best  conditions  under  which  to  bring 
about  the  reaction,  a  large  amount  of  data  had  to  be  accumulated 
from  accurate  measurements  of  the  factors  involved,  and  an  exhaus- 
tive search  made  for  a  proper  catalyzer.  It  was  found  that  the 
best  practical  conditions  were  a  temperature  of  450°,  a  pressure 
of  200  atmospheres,1  and  a  catalyzer  of  iron  which  contained  small 
quantities  of  other  substances  the  nature  of  which  is  a  trade  secret. 
To  fulfill  these  conditions  was  a  difficult  engineering  feat;  at  the 
high  temperature  required  hydrogen  passes  slowly  through  most 
metals,  and  their  strength  decreases.  To  construct  an  apparatus 
of  the  size  required  to  resist  such  a  high  pressure  was,  therefore, 
difficult.  By  using  steel  of  a  special  composition  and  by  lining 
the  apparatus  with  an  alloy  that  resisted  the  action  of  hydrogen 
at  the  temperature  used,  the  problem  was  finally  solved.  During 
the  course  of  the  work,  however,  explosions  of  great  violence 
occurred  and  great  steel  tubes,  which  contained  the  catalyzer, 
and  were  50  feet  long,  3  feet  in  diameter,  and  9  inches  thick  were 
blown  to  pieces. 

340.  The  production  of  the  nitrogen  and  the  hydrogen  for  the 
synthesis  presented  a  problem  which  was  solved  in  an  equally 
thorough  and  scientific  way.  When  air  is  passed  over  hot  coke 
the  nitrogen  is  unaffected  and  the  oxygen  converted  into  either 
carbon  dioxide  or  carbon  monoxide  or  a  mixture  of  the  two,  the 
temperature  and  other  conditions  determining  the  composition  of 
the  resulting  mixture  (200,  211) : 

C  +  O2  =  CO2 
CO2  +  C  +±  2CO 

If  the  oxides  are  removed  from  the  product,  the  reaction  furnishes  a 
source  of  nitrogen.  When  steam  is  passed  over  hot  coke,  hydrogen 
and  carbon  monoxide  and  carbon  dioxide  are  formed  (228) : 

H20  +  C  =  H2  +  CO 
CO  +  H2O  <±  CO2  +  H2 

1  It  has  been  reported  recently  that  Claude,  a  Frenchman,  has  constructed 
apparatus  in  which  the  reaction  can  be  carried  out  at  1000  atmospheres.  This 
change  would  increase  the  percentage  of  ammonia  in  the  gases  at  equilibrium. 


308  INORGANIC  CHEMISTRY  FOR  COLLEGES 

If  the  oxides  of  carbon  are  removed  in  this  case  the  reaction  fur- 
nishes a  source  of  hydrogen. 

By  passing  air  and  steam  together  over  hot  coke  the  resulting 
gas  is  a  mixture  of  nitrogen,  hydrogen,  and  the  two  oxides  of  carbon. 
The  proportions  of  the  several  constituents  are  determined  by  the 
relative  amounts  of  air  and  steam  used  and  the  temperature  at 
which  the  reaction  is  carried  out.  Since  carbon  monoxide  is 
difficult  to  remove  from  a  gas  on  account  of  its  lack  of  solubility 
in  common  liquids,  and  carbon  dioxide  can  be  readily  removed, 
it  is  necessary  to  run  the  reaction  under  the  conditions  which  yield 
the  highest  attainable  percentage  of  the  dioxide.  These  condi- 
tions can  be  determined  by  a  study  of  the  equilibria  in  the  revers- 
ible reactions  noted  above.  As  it  was  not  possible  to  convert 
all  the  carbon  monoxide  into  dioxide  in  a  single  operation,  the 
mixture  of  gases  first  obtained  was  mixed  with  a  fresh  supply  of 
steam  and  passed  over  a  catalyzer  which  increased  the  rate  of  the 
reaction  CO  +  H^O  <=±  C02  +  Eb.  The  catalyzer  in  this  case 
is  a  special  variety  of  iron  oxide.  The  reaction  just  given  is  a 
reversible  one,  so  the  carbon  monoxide  was  not  completely  removed. 
The  trace  that  remained  interfered  with  the  union  of  nitrogen  and 
hydrogen  and  was  dissolved  out  by  passing  the  gas  through  a 
solution  containing  a  cuprous  salt  (214).  To  remove  the  large 
amount  of  carbon  dioxide  in  the  gases  by  means  of  lime  was 
impracticable  on  account  of  the  expense,  and  water  was  used. 
The  solubility  of  a  gas  in  water  increases  rapidly  with  increased 
pressure  (Henry's  law,  185),  and  as  the  gas  had  to  be  compressed 
for  the  final  synthesis  it  did  not  add  to  the  cost  to  take  advantage 
of  the  increased  solubility  of  carbon  dioxide  under  these  condi- 
tions. By  washing  it  at  50  atmospheres  pressure  with  water, 
the  carbon  dioxide  was,  accordingly,  removed.  A  large  part  of 
the  energy  of  the  compressed  carbon  dioxide  was  utilized  by 
allowing  it  to  escape  from  the  water  solution  into  a  turbine  engine 
which  furnished  power. 

341.  To  sum  up,  the  Haber  process  for  making  synthetic 
ammonia  is  as  follows:  Nitrogen  and  hydrogen  are  obtained  by 
passing  air  and  steam  over  hot  coal  or  coke;  the  resulting  gases, 
mixed  with  more  steam,  are  passed  over  a  catalyzer  in  the  presence 
of  which  the  carbon  monoxide  present,  by  reacting  with  the  water, 
is  converted  into  carbon  dioxide.  The  gases  are  compressed  to 


AMMONIA  AND  ITS  DERIVATIVES  309 

50  atmospheres  and  brought  into  contact  with  water  which 
dissolves  out  the  carbon  dioxide.  They  are  next  passed  at  this 
pressure  through  a  solution  of  cuprous  formate  in  ammonia  which 
removes  the  trace  of  carbon  monoxide  remaining.  After  com- 
pression to  200  atmospheres  the  mixture  of  nitrogen  and  hydro- 
gen is  passed  over  a  catalyzer  and  unites  to  form  ammonia. 
Since  under  the  most  favorable  conditions  only  about  2  per  cent  of 
the  gases  react,  they  are  passed,  still  under  the  high  pressure, 
through  water  which  absorbs  the  ammonia,  and  are  then  returned 
to  the  catalyzer  chambers.  The  circulation  of  the  gases  is  con- 
tinued and  additional  nitrogen  and  hydrogen  admitted  at  the 
rate  at  which  they  unite.  The  water  containing  the  ammonia  is 
drawn  off  when  it  contains  enough  of  the  gas  to  saturate  it  at 
atmospheric  pressure  and  temperature.  The  ammonia  is  used 
as  such  or  converted  into  nitric  acid  in  a  way  to  be  described  later. 

The  Haber  plant  cost  about  $25,000,000  and  about  seventy-five 
chemists  and  physicists  are  engaged  in  the  control  of  the  process 
and  in  doing  the  research  work  required  for  its  development.  The 
production  of  nitric  acid  manufactured  from  the  ammonia  made 
in  this  way  was  approximately  100,000  tons  per  year. 

The  synthesis  of  ammonia  by  the  Haber  process  has  been 
described  in  some  detail  because  it  is  an  excellent  illustration  of 
how  a  solution  of  an  important  industrial  chemical  problem  has 
been  reached  through  the  application  of  physical  chemistry.  It 
emphasizes  the  importance  of  the  study  of  chemical  equilibrium 
and  the  factors  which  govern  it,  all  of  which  are  susceptible  of 
mathematical  analysis.  We  have  seen  how  van't  HofFs  law, 
Henry's  law,  and  the  principle  of  Le  Chatelier,  have  been  applied 
to  the  real  service  of  mankind.  The  development  of  industrial 
chemistry  in  all  its  phases  is  rapidly  taking  place  as  the  result  of 
the  application  of  the  fundamental  principles  underlying  the 
science,  which  are  considered  fully  in  physical  or  theoretical 
chemistry. 

342.  The  Cyanamide  Process  for  Ammonia. — Where  elec- 
trical power  is  cheap  this  process  for  the  production  of  ammonia 
can  be  advantageously  used.  Calcium  cyanamide  decomposes 
slowly  with  water  and  yields  ammonia,  and  for  this  reason  is 
used  itself  as  a  fertilizer,  as  it  furnishes  a  cheap  source  of  this 
important  plant  food.  The  compound  has  been  made  in  large 


310  INORGANIC  CHEMISTRY  FOR  COLLEGES 

quantities  in  Norway  and  in  the  United  States  where  water  power 
for  generating  electricity  is  available.  In  Germany  brown  coal 
and  lignite  are  abundant,  and  furnish  a  cheap  source  of  power. 
When  nitrogen  is  passed  over  calcium  carbide  (217)  at  the  tem- 
perature of  an  electric  furnace,  the  gas  is  absorbed  and  calcium 
cyanamide  is  formed : 

CaC2  +  N2  =  CaCN2  +  C 

The  compound  decomposes  slowly  with  water  with  the  formation 
of  ammonia: 

CaCN2  +  3H2O  =  CaCO3  +  2NH3 

If  steam  is  used  the  decomposition  is  rapid  and  by  condensing  the 
vapor  a  solution  of  ammonia  is  obtained. 

In  calcium  cyanamide  the  calcium  is  in  combination  with  the 
nitrogen.  It  will  be  recalled  (318)  that  compounds  of  metals  with 
nitrogen  are  decomposed  by  water,  and  that  the  hydrogen  of  the 
latter  unites  with  the  nitrogen  to  form  ammonia,  and  the  oxygen 
with  the  metal  to  form  an  oxide.  The  reaction  of  calcium  cyana- 
mide with  water  takes  place  in  an  analogous  way;  in  this  case, 
however,  the  carbon  in  the  compound  is  converted  into  carbon 
dioxide  which  reacts  with  the  calcium  oxide  and  forms  calcium 
carbonate. 

343.  The  formation  of  ammonia  when  animal  matter  undergoes 
decomposition  has  been  emphasized.  An  example  of  this  is  the 
decomposition  of  urine,  which  results  in  the  conversion  of  the  urea 
present  into  ammonia  and  carbon  dioxide,  through  the  agency  of  a 
micro-organism  present  in  the  air. 

Ammonia  is  formed  when  the  proteins  and  other  nitrogenous 
materials  in  plants  and  animals  are  heated  with  concentrated  sul- 
phuric acid.  The  acid  oxidizes  these  compounds,  and  the  carbon 
they  contain  is  changed  into  carbon  dioxide,  the  hydrogen  into 
water,  and  the  nitrogen  into  ammonia.  The  latter  unites  with  the 
excess  of  sulphuric  acid  to  form  ammonium  sulphate.  Advantage 
is  taken  of  this  fact  in  analyzing  quantitatively  food  products  for 
nitrogen.  Proteins  are  essentials  in  food  and  their  determination 
in  the  valuation  of  a  food  product  is,  therefore,  important.  The 
product  to  be  analyzed  is  heated  with  concentrated  sulphuric 
acid,  and  when  the  oxidation  is  complete  the  solution  is  made 


AMMONIA  AND  ITS  DERIVATIVES  311 

alkaline  with  sodium  hydroxide  and  heated.  The  ammonia  set 
free  distills  over  with  water;  its  amount  is  determined  by  neu- 
tralizing it  with  an  acid  of  known  concentration.  The  procedure 
outlined  is  used  in  what  is  called  the  Kjeldahl  method  for  the  deter- 
mination of  nitrogen.  The  oxidation  with  sulphuric  acid  is  cata- 
lyzed by  mercury  salts. 

344.  Physical  Properties  of  Ammonia. — Ammonia  is  a  colorless 
gas  which  is  lighter  than  air  (sp.  gr.  0.5971).     Its  critical  tempera- 
ture is  131°  and  it  can,  therefore,  be  condensed  to  a  liquid  at 
ordinary  temperatures.     Liquid  ammonia  boils  at    —33.5°  and 
solidifies  to  a  white  crystalline  solid  at  -77°.     Its  molecular  heat 
of  vaporization  is  5700  calories,  the  value  being  greater  than  that 
of  any  other  liquid  except  water.     Ammonia  is  very  soluble  in 
water;    1  volume  of  water  at  0°  dissolves  1300  volumes  of  the 
gas,  and  at   16°,  783  volumes.     The  concentrated  ammonia  of 
commerce  contains  about  28  per  cent  by  weight  of  the  gas  and 
has  the  specific  gravity  0.9. 

Liquid  ammonia  is  an  excellent  solvent  for  both  inorganic  and 
organic  substances.  It  dissolves  sodium  and  potassium  and  other 
substances  which  are  insoluble  in  or  react  with  water.  When  salts 
dissolve  in  liquid  ammonia  they  conduct  the  electric  current  and 
show  the  characteristic  behavior  exhibited  when  ionized  in  water. 
The  liquid  is  transported  in  iron  cylinders  and  is  used  in  ice- 
machines  (180).  The  pressure  in  a  cylinder  of  liquid  ammonia  at 
20°  (68°  F.)  is  about  8  atmospheres,  and  at  30°  (86°  F.)  about  11.5 
atmospheres. 

345.  Chemical  Behavior  of  Ammonia. — The  heat  of  formation 
of  ammonia  is  small  compared  with  that  of  water  and  we  would 
expect,  therefore,  to  find  it  less  stable  toward  heat.     When  1  gram- 
molecule  (17  grams)  of  ammonia  is  formed  from  nitrogen  and 
hydrogen  12,200  calories  are  set  free;   the  production  of  a  corre- 
sponding amount  of  water- vapor  (18  grams)  yields  58,100  calories. 
The   gas   is   almost   completely   decomposed   into   nitrogen   and 
hydrogen  at  700°. 

The  change  in  volume  which  occurs  when  ammonia  decomposes 
can  be  determined  by  passing  electric  sparks  through  a  measured 
volume  of  the  gas  collected  in  a  glass  tube  over  mercury.  When 
the  sparking  no  longer  produces  a  change  it  will  be  found  that  the 
volume  of  the  gases  produced  is  twice  that  of  the  ammonia  taken. 


312  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  relative  proportions  of  the  hydrogen  and  nitrogen  in  the 
resulting  gases  can  be  determined  by  passing  them  over  hot  copper 
oxide,  which  converts  the  hydrogen  into  water.  The  results 
show  that  the  volume  of  hydrogen  is  three  times  that  of  the 
nitrogen.  The  equation  for  the  reaction  is 

2NH3  <=±  N2  +  3H2 

When  ammonia  is  heated  with  the  more  active  metals,  such  as 
magnesium,  nitrides  are  formed : 

3Mg  +  2NH3  =  Mg3N2  +  3H2 

At  lower  temperatures  with  sodium,  only  a  part  of  the  hydrogen 
is  set  free  and  sodium  amide  results : 

2Na  +  2NH3  -  2NaNH2  +  H2 

Ammonia  forms  addition-products  with  certain  salts,  which 
resemble  in  composition  the  hydrates  of  these  salts;  for  example, 
compounds  of  the  composition  CaCl2,6NH3  and  CaCl2,6H2O, 
and  CuSO4,5NH3  and  CuSO4,5H20  are  known.  It  is  evident 
that  anhydrous  salts  which  unite  with  ammonia  cannot  be  used 
for  drying  ammonia  gas  containing  water-vapor.  The  solubility 
in  water  of  the  compounds  of  salts  with  ammonia,  is  often  quite 
different  from  the  solubility  of  these  salts  themselves.  For 
example,  silver  chloride,  AgCl,  is  almost  insoluble  in  water,  whereas 
its  compound  with  ammonia,  AgCl,2NH3,  is  readily  soluble  in 
water.  It  is  for  this  reason  that  precipitated  silver  chloride  dis- 
solves in  a  solution  of  ammonia.  The  change  in  solubility  when 
many  inorganic  substances  are  treated  with  ammonia  is  made  use 
of  in  qualitative  analysis. 

Ammonia  does  not  burn  when  a  stream  of  the  gas  issuing  from  a 
tube  is  brought  into  contact  with  a  flame.  If  surrounded  by  an 
atmosphere  of  oxygen,  however,  it  burns  with  a  yellow  flame  on 
ignition.  This  can  be  easily  demonstrated  by  passing  oxygen  to 
the  bottom  of  a  vessel  containing  a  concentrated  aqueous  solution 
of  the  gas;  as  the  oxygen  rises  through  the  liquid  it  carries  along 
with  it  some  ammonia,  and  the  gases  when  they  leave  the  solution 
can  be  ignited;  they  burn  with  a  yellow  flame.  The  product 


AMMONIA  AND  ITS  DERIVATIVES  313 

formed  depends  upon  the  proportion  of  oxygen;  when  an  excess 
of  ammonia  is  present  nitrogen  is  formed  : 

4NH3  +  3O2  =  2N2  -f  6H2O 

With  more  oxygen,  oxides  of  nitrogen  are  produced. 

By  burning  ammonia  with  air  in  the  presence  of  a  catalyzer 
nitric  oxide  can  be  formed  : 

4NH3  +  5O2  =  4NO  +  6H2O 

We  shall  see  later  that  nitric  oxide  unites  with  oxygen  from  the  air 
in  the  presence  of  water  to  form  nitric  acid.  The  reaction  is  the 
basis  of  an  important  new  chemical  industry  which  will  be  described 
in  the  next  chapter. 

An  aqueous  solution  of  ammonia  shows  all  the  characteristic 
properties  of  a  solution  of  a  base.  For  this  reason  it  was  believed 
for  many  years  that  ammonia  reacted  with  water  to  form  a  com- 
pound of  the  composition  NH-iOH,  although  no  one  had  suc- 
ceeded in  isolating  it.  Very  recently  this  compound,  ammonium 
hydroxide,  has  been  isolated  at  low  temperatures;  it  forms 
crystals  which  melt  at  —79°. 

346.  The  reaction  between  ammonia  and  water  is  a  reversible 
one: 

NH3  +  H2O  ^>  NH4OH 

It  has  been  shown  that  in  a  molar  solution  (17  grams  of  NH3  in 
1  liter  of  solution)  at  20°  about  30  per  cent  of  the  ammonia  is  in 
combination  with  water  as  ammonium  hydroxide.  As  the  tem- 
perature rises  the  decomposition  indicated  by  the  equation  when 
read  from  right  to  left  takes  place  rapidly,  and  at  the  boiling- 
point  of  water  it  is  practically  complete.  Water  can  be  com- 
pletely freed  from  ammonia  by  boiling  it. 

Ammonium  hydroxide  is  a  comparatively  weak  base,  that  is,  the 
extent  to  which  it  ionizes  in  water  is  much  less  than  is  the  case 
with  sodium  hydroxide;  it  ionizes  as  indicated  below: 

NH4OH  =  NH4+ 


In  a  molar  solution  of  sodium  hydroxide  72  per  cent  of  the  base  is 
in  the  form  of  ions,  whereas  in  a  molar  solution  of  ammonia  only 
0.4  per  cent  is  ionized.  In  solutions  of  one-tenth  this  concen- 


314  INORGANIC  CHEMISTRY  FOR  COLLEGES 

tration  the  ionization  values  are  91  per  cent  and  1.3  per  cent,  respect- 
ively. The  caustic  properties  of  a  base  are  due  to  the  hydroxyl 
ions  which  it  produces.  Sodium  hydroxide,  for  example,  decomposes 
and  dissolves  protein  material,  and  for  this  reason  when  a  solution 
of  it  is  allowed  to  stay  in  contact  with  the  flesh  a  painful  "  burn  " 
results;  with  ammonium  hydroxide,  however,  this  does  not  occur 
because  the  concentration  of  the  hydroxyl  ions  in  the  solution  is 
so  small. 

347.  Uses  of  Ammonia. — Ammonium  hydroxide  is  used  in  the 
household  for  softening  water  for  laundry  purposes.     We  shall  see 
later  that  this  is  brought  about  by  causing  the  precipitation  of  the 
compounds  present  in  the  water  which  interfere  with  the  action  of 
soap.     Ammonium  hydroxide  also  acts  as  a  cleansing  agent  by 
converting  the  oils  and  grease  present  into  exceedingly  fine  glob- 
ules which  do  not  adhere  to  one  another  and  so  can  be  removed  by 
water.     Ammonia  is  of  value  in  cleaning  certain  metals  from  rusts 
and  deposits  caused  by  the  presence  of  sulphur  dioxide  and  other 
gases  in  the  air;  the  use  for  this  purpose  is  based  on  the  fact  that 
ammonia  reacts  with  the  salts  present  by  direct  addition  and 
converts  them  into  compounds  soluble  in  water.     Ammonia  is 
taken  internally  as  a  medicine,  and  is  a  heart  stimulant.     It  is  the 
gas  liberated  by  "  smelling  salts."     In  larger  quantities  it  is  an 
active  poison. 

348.  Ammonium  Salts. — Ammonia  unites  directly  with  acids 
and  forms  ammonium  salts;    ammonium  chloride  is  formed  when 
ammonia  and  hydrogen  chloride  are  brought  into  contact  (140). 
This  experiment  was  first  performed  by  Priestley.     He  had  iso- 
lated ammonia  gas,  which  he  called  alkaline  air,  and  hydrogen 
chloride,  which  he  called  acid  air;   and  in  an  endeavor  to  get  a 
neutral  air  he  brought  the  two  together  and,  as  a  result,  discovered 
the  composition  of  ammonium  chloride,  which  was  up  to  that  time 
unknown. 

Ammonium  salts  are  also  formed  by  neutralizing  ammonium 
hydroxide  with  solutions  of  acids,  for  example,  ammonium  sul- 
phate can  be  formed  in  this  way: 

2NH4OH  +  H2SO4  =  (NH4)2SO4  +  2H2O 

The  salts  derived  from  ammonia  are  called  ammonium  salts  because 
they  resemble  in  properties  the  salts  of  metals,  the  names  of  which 


AMMONIA  AND  ITS  DERIVATIVES  315 

in  most  cases  terminate  in  ium,  for  example,  sodium,  calcium, 
aluminium.  All  ammonium  salts  contain  the  combination  of 
atoms  represented  by  the  symbols  NEU;  this  is  called  a  radical 
and  passes  unchanged  from  one  ammonium  compound  to  another. 
It  plays  the  same  part  in  ammonium  salts  that  metallic  atoms  play 
in  their  salts;  it  passes  into  the  form  of  an  ion  when  ammonium 
salts  are  dissolved  in  water,  and  has  the  valence  1.  When  attempts 
were  made  to  isolate  it,  decomposition  into  ammonia  and 
hydrogen  took  place.  But  some  of  these  experiments  lead  to  the 
conclusion  that  the  radical  possesses  the  properties  of  metals. 
When  an  electric  current  is  passed  through  a  solution  of  sodium 
chloride  and  the  cathode  is  mercury,  the  metallic  sodium  liberated 
dissolves  in  the  mercury  and  forms  a  compound  with  it,  which  is 
hard  and  shows  the  properties  of  an  alloy.1  When  a  solution  of 
ammonium  chloride  is  electrolyzed  under  the  same  conditions 
below  0°  a  solid  ammonium  amalgam  is  obtained  which  resembles 
the  corresponding  product  containing  sodium.  If  the  temperature 
is  allowed  to  rise,  the  amalgam  softens  and  swells  up  to  a  spongy 
mass  as  the  result  of  the  decomposition  of  the  ammonium  into 
ammonia  and  hydrogen.  These  facts  indicate  that  the  radical 
ammonium  exhibits  the  properties  of  metals  in  alloys  when  it  is 
in  combination  with  another  metal;  and  it  has  already  been  shown 
that  the  radical  plays  the  part  of  a  metal  when  it  is  in  combination 
with  acid  radicals. 

A  large  number  of  ammonium  salts  are  known  and  many  of 
them  have  interesting  applications.  Ammonium  chloride,  sal 
ammoniac,  is  obtained  in  impure  condition  by  neutralizing  the 
gases  given  off  when  coal  is  distilled  in  making  coke  or  illuminating 
gas.  It  is  readily  purified  by  sublimation  as  its  vapor  pressure  is 
equal  to  the  pressure  of  the  atmosphere  at  about  338°.  Large 
quantities  of  the  salt  are  used  in  the  manufacture  of  dry  cells,  as 
a  source  of  ammonia,  and  in  chemical  laboratories.  It  is  used  as  a 
flux  in  soldering,  as  it  serves  to  remove  the  coating  of  oxides  formed 
on  the  solder  and  the  metal  and  thus  makes  it  possible  for  the  two 

1  Alloys  are  usually  produced  by  melting  two  metals  together;  brass  is 
prepared  from  copper  and  zinc  in  this  way.  Alloys  often  contain  compounds 
of  definite  composition  formed  as  the  result  of  the  union  of  the  metals  from 
which  they  are  prepared.  The  alloys  containing  mercury  are  called  amal- 
gams; sodium  amalgam,  for  example,  contains  a  compound  of  the  formula 
NaHga. 


316  INORGANIC  CHEMISTRY  FOR  COLLEGES 

to  alloy.  The  ammonium  chloride  dissolves  the  oxides  because 
when  it  is  heated  it  dissociates  into  ammonia  and  hydrogen  chloride, 
and  the  latter  converts  the  oxides  into  chlorides,  which  melt  and 
run  off. 

Crude  ammonium  sulphate,  (NH4)2SO4,  is  obtained  from  the 
gas  works  and  is  used  chiefly  as  an  ingredient  of  fertilizers. 

349.  During  the  recent  war  the  supply  of  sulphuric  acid  in 
Germany  was  limited  on  account  of  the  difficulty  of  obtaining 
sulphur  and  sulphides.     Large  quantities  of  ammonium  sulphate 
were  manufactured  from  calcium  sulphate  directly  without  first 
preparing  sulphuric  acid.     Finely  ground  gypsum,  the  form  in 
which  calcium  sulphate  occurs  in  nature,  was  stirred  in  water 
and  treated  with  ammonia  and  carbon  dioxide.     The  ammonium 
carbonate  produced  reacted  with  the  sulphate  and  formed  am- 
monium sulphate  and  calcium  carbonate: 

CaSO4+  (NH4)2CO3  =  (NH4)2S04  +  CaCO3 

After  filtrating  from  the  insoluble  carbonate,  the  solution  on  evap- 
oration yielded  ammonium  sulphate.  This  process  avoids  the 
manufacture  of  sulphuric  acid  and  is  of  economic  significance. 

350.  Ammonium  carbonate,  (NH^COa,  is  formed  by  passing 
carbon  dioxide  through  a  solution  of  ammonia,  if  precautions 
are  taken  to  have  an  excess  of  the  latter  present.     On  evaporation 
the  normal  carbonate  is  obtained,  but  it  is  quite  unstable  and  when 
left  in  the  air  soon  loses  ammonia  an  1  is  changed  into  the  bicar- 
bonate,  NH4HC03.       The  latter  can  be  made  by  completely 
saturating  a  solution  of  ammonia  with  carbon  dioxide  and  subse- 
quent evaporation.     The  salt  is  obtained  as  white  crystals,  which 
slowly  give  off  ammonia.     This  property  leads  to  its  use  in  "  smell- 
ing salts."     Ammonium  bicarbonate  is  moistened  with  a  solution 
containing  a  perfume  of  pleasant  odor;    when  the  bottle  con- 
taining the  salts  is  opened  the  ammonia  which  has  accumulated  is 
given  off  in  the  desired  concentration.     The  uses  to  which  ammo- 
nium sulphide  and  polysulphide  (282,  311)  and  ammonium  per- 
sulphate are  put  have  been  mentioned;  other  salts  will  be  described 
later. 

351.  General  Properties  of  Ammonium  Salts. — All  these  salts 
are  decomposed  when  heated,  but  the  behavior  of  any  particular 
one  is  determined  by  the  acid  radical  present.     If  the  acid  from 


AMMONIA  AND  ITS  DERIVATIVES  317 

which  the  salt  is  derived  is  volatile,  the  salt  sublimes  and  if  pure 
leaves  no  residue;  if  it  is  not  volatile,  the  acid  or  its  anhydride  is 
left  when  the  salt  is  ignited.  It  is  for  this  reason  that  ammonium 
hydroxide  is  used  frequently  in  analytical  chemistry  when  it  is 
necessary  in  the  process  to  neutralize  an  acid ;  the  solution  obtained 
can  be  evaporated  to  dryness  and  heated,  and  thus  freed  from  the 
base  used  during  the  course  of  an  analysis,  in  which  the  presence 
of  sodium  or  potassium  would  interfere  with  obtaining  the  results 
desired. 

All  ammonium  salts  with  the  exception  of  a  few  so-called 
double  salts,  like  the  one  of  the  composition  PtCL^NEUCl,  are 
readily  soluble  in  water.  The  fact  that  all  ammonium  salts  are 
decomposed  in  aqueous  solution  by  a  base  is  utilized  in  the  general 
test  for  these  salts.  The  test  is  carried  out  by  treating  a  solution  of 
the  material  under  examination  with  a  solution  of  sodium  hydrox- 
ide and  noting  the  odor;  if  ammonia  is  not  present  the  solution  is 
heated  and  the  odor  again  noted.  All  ammonium  salts  yield 
ammonia  under  these  conditions. 

352.  Other  compounds  of  Nitrogen  and  Hydrogen. — It  has 
been  pointed  out  that  nitric  acid  can  be  reduced  by  nascent 
hydrogen  to  ammonia.  When  the  most  active  reducing  agents 
are  used  with  the  acid  or  with  the  oxides  of  nitrogen,  the  final 
reduction-product  in  all  cases  is  ammonia.  When  the  less  active 
reducing  agents  are  used,  however,  under  the  proper  conditions 
other  reduction-products  are  formed.  For  example,  when  nitric 
oxide,  NO,  is  reduced  by  combining  the  action  of  potassium  sul- 
phite with  that  of  the  nascent  hydrogen  generated  by  the  action  of 
water  on  sodium  amalgam  (note,  348),  a  compound  of  the  formula 
N2H4,  called  hydrazine  y  is  formed.  In  this  compound  the  reduction 
of  the  nitrogen  has  not  proceeded  as  far  as  is  the  case  with  ammo- 
nia; two  nitrogen  atoms  in  the  form  of  ammonia  are  in  combi- 
nation with  six  hydrogen  atoms,  whereas  in  hydrazine  two  nitrogen 
atoms  are  united  with  four  atoms  only.  If  hydrazine  is  treated 
with  hydrogen  generated  by  a  metal  and  an  acid  it  is  reduced 
further  to  ammonia. 

Hydrazine  is  a  colorless  liquid  which  boils  at  113°.  It  unites 
with  water  to  form  a  compound  which  unlike  ammonium  hydrox- 
ide is  stable.  Hydrazine  and  its  compound  with  water  resemble 
ammonia  and  ammonium  hydroxide  in  chemical  properties. 


318  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Since  it  contains  two  atoms  of  nitrogen  in  the  molecule  hydrazine 
can  form  salts  by  uniting  with  either  one  or  two  molecules  of  a 
monobasic  acid.  Hydrazine  is  readily  oxidized  and  is,  therefore, 
an  active  reducing  agent. 

353.  When  dilute  nitric  acid  is  reduced  at  low  temperatures 
with  nascent  hydrogen  generated  by  the  action  of  tin  on  the  acid, 
all  but  one  of  the  oxygen  atoms  are  removed  and  three  hydrogen 
atoms  are  added.    The  resulting  product  is  called  hydroxylamine, 
and  has  the  formula  NH2OH.     It  is  a  white  solid  which  melts  at 
33°.     The  compound  forms  salts  by  direct  addition  with  acids, 
and  its  aqueous  solution  shows  the  properties  of  a  weak  base. 
Like  hydrazine,  hydroxylamine  is  an  active  reducing  agent. 

354.  Hydrazoic  acid,  HN3,  is  a  compound  of  peculiar  interest 
on  account  of  the  fact  that  it  is  so  unstable.     It  can  be  formed 
by  the  oxidation  of  hydrazine  with  n;trous  acid: 

N2H4  +  HNO2  =  HN3  +  2H2O 

The  sodium  salt  of  the  acid  is  more  conveniently  prepared  by 
treating  sodium  amide  with  nitrous  oxide: 

NaNH2  +  N2O  =  NaN3  +  H2O 

Since  the  acid  is  volatile  with  steam  it  can  be  obtained  by  distilling 
the  sodium  salt  with  dilute  sulphuric  acid.  Hydrazoic  acid,  which 
is  also  called  hydronitric  acid,  is  a  colorless  liquid  which  possesses  a 
strong,  disagreeable  odor;  it  boils  at  37°  and  explodes  with  great 
violence  when  heated.  The  reaction  is  accompanied  by  the  evo- 
lution of  a  large  quantity  of  heat: 

2HN3  =  3N2  +H2  +  2  X  61,600  cals. 

The  salts  of  the  acid,  which  are  called  azides,  are,  in  the  case  of  the 
heavy  metals,  explosive  compounds.  Lead  azide  has  recently 
come  into  use  in  the  manufacture  of  caps  for  high  explosive  shell. 

EXERCISES 

1.  Ammonia  was  passed  over  hot  copper  oxide  and  the  nitrogen  formed 
was  determined:  2NH3+3CuO=  N2 -f  3H2O  +  3Cu.  (a)  Calculate  the 
percentage  composition  of  NH3  from  the  following  results:  Weight  CuO 
before  experiment  6.450  grams;  weight  CuO  after  experiment  6.258;  volume 
of  N2  at  0°  and  760  mm.  89.6  c.c.  (6)  Calculate  the  percentage  composition 
of  NH3  from  its  formula, 


AMMONIA  AND  ITS  DERIVATIVES  319 

2.  The  substances  having  the  following  formulas  are  commonly  used  in 
the  laboratory  as  drying  agents:   CaCl2,  KOH,  CuSO4,  H2SO4,  P2O6.     Which 
can  and  which  cannot  be  used  to  dry  ammonia?     Give  a  reason  in  each  case. 

3.  (a)  Calculate  from  its  formula  the  weight  of  1  liter  of  NH3.    (6)  Cal- 
culate the  specific  gravity  of  NH3  compared  with  air,  1  liter  of  which  weighs 
1.298  grams,     (c)  Could  NH3  be  used  in  balloons?     What  evident  disadvan- 
tage would  the  gas  possess  for  this  purpose? 

4.  How  could  you  determine  the  volume  of  NH3  in  a  mixture  of  the  latter 
and  air  contained  in  a  flask? 

5.  Write  equations  for  three  reactions  by  which  NH3  may  be  produced 
from  three  different  salts,  using  a  different  base  in  each  case. 

6.  Compare  the  action  at  high  temperatures  of  H^O  and  NH3  on  metals. 

7.  State  three  ways  in  which  you  could  show  that  a  solution  of  NH3  in 
water  contains  OH  ions. 

8.  What  would  you  expect  to  be  formed  when  an  electric  current  is  passed 
through  a  solution  of  (NH4)2SO4  in  water? 

9.  How   could  you  distinguish  from    one    another    the    following:     (a) 
(NH4)2SO4,     (6)  NH4C1,    (c)  NH4NO3,    (d)  NaCl,    (e)  Na2SO4,    (/)  NaOH, 
(0)  NH4HC03,  (A)  (NH4)2S? 

10.  Name  the  compounds  and  ions  present  in  a  solution  of  NH3  in  water 
and  represent  by  equations  the  equilibrium  that  exists  between  them.  Show 
how  the  equilibrium  changes  in  each  reaction  when  the  solution  is  heated 
to  boiling, 


CHAPTER  XXIII 
NITRIC  ACID,  NITROUS  ACID,  AND  THE  OXIDES  OF  NITROGEN 

355.  Nitric  acid,  nitrous  acid  and  the  oxides  of  nitrogen  are  of 
particular  interest  and  significance;  their  chemical  properties 
illustrate  some  of  the  most  fundamental  principles  underlying  the 
behavior  of  molecules,  and  the  energy  changes  involved  in  their 
transformations  furnish  an  opportunity  to  grasp  clearly  the  sig- 
nificance of  the  energy  factor  in  chemical  change.  We  have  studied 
in  some  detail  carbon  monoxide,  CO,  and  carbon  dioxide,  CO2, 
and  have  learned  that  these  compounds  can  be  readily  formed,  and 
that  in  their  production  large  quantities  of  heat  are  evolved — the 
reactions  are  exothermic.  It  will  be  of  interest  to  study  now  the 
nitrogen  compounds  of  analogous  composition,  nitric  oxide,  NO, 
and  nitrogen  dioxide,  NO2,  because  these  compounds  are  endo- 
thermic — that  is,  heat  is  absorbed  when  they  are  formed. 

The  tendency  in  nature  is  to  render  energy  unavailable — a  stone 
falls  to  the  ground  and  the  power  it  had  to  do  work  due  to  its 
position  is  lost,  and  the  energy  passes  into  heat,  only  a  part  of  which 
can  be  utilized.  Reactions  which  continue  to  take  place  when  once 
started  are,  in  most  cases,  those  which  give  off  energy.  Carbon 
is  a  storehouse  of  chemical  energy  that  can  be  utilized  to  do  the 
mechanical  work  of  the  world;  and  the  product  formed  when  it 
burns  freely  is  comparatively  inert;  millions  of  tons  of  carbon 
dioxide — the  product  of  burning  coal — are  allowed  to  escape, 
because  the  compound  contains  such  a  small  amount  of  available 
energy.  Nitrogen,  on  the  other  hand,  unites  with  oxygen  only 
as  long  as  we  continue  to  supply  a  large  amount  of  energy;  and 
the  resulting  oxide  can  be  used  as  a  source  of  energy.  It  is  an 
endothermic  compound,  and  as  it  is  an  excellent  example  of  the 
class  to  which  it  belongs  its  study  will  be  of  great  importance. 

The  tremendous,  destructive  force  shown  by  gunpowder,  dynam- 
ite, smokeless  powder,  and  the  other  high  explosives  used  in  warfare 

320 


NITRIC  ACID,  NITROUS  ACID,  OXIDES  OF  NITROGEN  321 

and  for  blasting  and  other  industrial  purposes,  is  traceable  to  the 
fact  that  when  a  nitrogen  atom  unites  with  oxygen  atoms  the 
energy  furnished  to  effect  the  union  is  stored  up  and  becomes 
available  when  the  substance  containing  these  atoms  reacts  with 
other  compounds.  Since  in  all  chemical  changes  matter  and  energy 
are  involved,  it  is  important  not  to  lose  sight  of  the  energy  factor 
when  considering  the  matter,  which,  in  most  cases,  more  often 
affects  our  senses.  In  all  industrial  problems  involving  the  prep- 
aration of  chemical  compounds,  the  skilled  chemical  engineer  pays 
particular  attention  to  the  energy  involved,  since  it  is  often  an 
important  factor  in  the  cost  of  the  manufactured  product. 

NITRIC   ACID 

356.  Nitric  acid,  HNOs,  does  not  occur  in  nature  in  the  free 
condition  except  in  traces  in  the  air  after  a  thunderstorm  (324). 
It  is  active  chemically  and  if  set  free  would  soon  find  something 
with  which  to  react.  It  occurs  as  nitrates  in  certain  parts  of  the 
earth.  Potassium  nitrate  is  found  in  the  soil  near  cities  in  certain 
Oriental  countries,  having  been  produced  as  the  result  of  the  action 
of  nitrifying  bacteria  on  animal  refuse.  Such  soil  was  the  source 
of  Bengal  saltpeter.  Large  deposits  of  sodium  nitrate  are  found 
in  Chile,  and  these  are  now  the  chief  source  of  the  combined 
nitrogen  used  in  the  world.  The  deposits  are  in  an  arid  region  of 
North  Chile  and  cover  about  500  square  miles;  they  are  from  2  to 
10  feet  thick  and  contain  up  to  60  per  cent  of  sodium  nitrate. 
Guano  (315),  formerly  much  used  as  a  fertilizer,  contains  a  large 
percentage  of  nitrates  which  have  been  produced  as  the  result  of 
bacterial  action  on  the  organic  nitrogen  compounds  present  in 
the  excreta  of  birds,  from  which  the  guano  was  formed. 

Nitrates  which  are  present  in  the  soil  are  used  by  the  plant 
in  building  up  the  organic  nitrogen  compounds  produced  in  their 
growth.  If  the  material  grown  is  taken  away,  in  a  short  time  the 
soil  becomes  depleted  and  it  is  necessary  to  renew  the  nitrates 
through  the  application  of  fertilizers.  Nitrates  are  found  in 
small  and  varying  amounts  in  water,  and  their  presence  and 
quantity  serve,  as  we  shall  see,  to  indicate  to  the  chemists  the 
history  of  the  water  as  regards  pollution. 

Nitric  acid  was  known  in  ancient  times,  and  was  called  aqua 


322  INORGANIC  CHEMISTRY  FOR  COLLEGES 

fortis  on  account  of  the  fact  that  it  was  an  excellent  solvent  for 
substances  which  were  not  affected  by  less  powerful  reagents. 
Lavoisier  (1776)  first  showed  that  it  contained  hydrogen,  nitrogen, 
and  oxygen,  and  Gay-Lussac  later  determined  its  exact  composi- 
tion. 

357.  Preparation  of  Nitric  Acid. — In  preparing  nitric  acid 
from  its  salts  we  make  use  of  the  important  fact  that  double 
decompositions  proceed  to  completion  provided  one  of  the  products 
of  the  reaction  is  a  gas.  Nitric  acid  is  a  liquid  at  the  ordinary 
temperature  but  boils  at  86°;  as  a  consequence,  if  a  nitrate  is  mixed 
with  sulphuric  acid  and  heated  to  this  temperature  the  conditions 
are  those  which  lead  to  double  decomposition.  The  equation  for 
the  reaction  used  in  the  technical  preparation  of  nitric  acid  is  as 
follows : 

NaNO3  +  H2SO4  =  NaHSO4  +  HNO3 

The  proportions  indicated  are  used  rather  than  those  which  lead 
to  the  formation  of  neutral  sodium  sulphate: 

2NaNO3  +  H2SO4  =  Na2SO4  +  2HNO3 

This  second  reaction  appears  to  be  the  more  economical  one 
because  1  molecule  of  sulphuric  acid  causes  the  liberation  of  2 
molecules  of  nitric  acid,  whereas  in  the  first  reaction  but  half  that 
amount  would  be  obtained;  the  cost  of  the  sulphuric  acid  used  in 
the  process  based  on  the  second  equation  would  be,  thus,  one-half 
the  cost  of  the  acid  used  in  the  first  one.  It  will  be  recalled  that 
when  a  reaction  takes  place  according  to  the  second  equation  it 
proceeds  in  steps;  that  represented  by  the  first  equation  first 
takes  place,  and  then  the  acid  sulphate  reacts  with  another 
molecule  of  nitrate  as  follows: 

NaNO3  +  NaHSO4  =  Na2SO4  +  HNO3 

The  temperature  at  which  this  double  decomposition  takes  place 
is  much  higher  than  that  required  for  the  first  step;  and  as  this 
temperature  is  above  that  at  which  nitric  acid  partially  decomposes 
the  reaction  cannot  be  used. 

In  manufacturing  nitric  acid,  sodium  nitrate  is  heated  in 
cast-iron  retorts  with  sulphuric  acid  having  a  specific  gravity  of 
1.71  to  1.83.  depending  on  the  strength  of  acid  desired.  The 


NITRIC  ACID,  NITROUS  ACID,  OXIDES  OF  NITROGEN     323 

escaping  vapor  is  condensed  in  glass,  stone-ware,  or  aluminium 
pipes  surrounded  by  water.  To  free  the  acid  from  the  small 
amounts  of  oxides  of  nitrogen  formed  in  the  distillation,  which  give 
it  a  yellow  color,  air  is  blown  through  the  acid. 

If  concentrated  nitric  acid  and  concentrated  sulphuric  acid  are 
mixed  and  heated  cautiously,  nitric  acid  free  from  water  distills 
over.  It  contains  nitrogen  dioxide  produced  as  the  result  of  the 
decomposition  of  a  part  of  the  nitric  acid  and,  as  a  result,  has  a 
brown  color.  The  acid  fumes  in  the  air  and  is  called,  therefore, 
fuming  nitric  acid ;  it  has  the  specific  gravity  1 .6  and  is  a  powerful 
oxidizing  agent. 

358.  Synthetic  Nitric  Acid. — The  fact  has  been  known  for  a 
long  time  that  when  an  electric  spark  is  passed  through  a  mixture 
of  nitrogen  and  oxygen  the  elements  unite  and,  if  water  is  present, 
nitric  acid  is  formed.  Although  Cavendish  carried  out  this 
experiment  in  1785  it  has  only  been  in  recent  years  that  it  has  been 
possible  to  apply  the  reactions  involved  to  the  manufacture  of 
nitric  acid.  Science  had  to  advance  as  the  result  of  the  discovery 
of  the  fundamental  laws  underlying  chemical  equilibrium,  before 
it  was  possible  to  develop  a  great  chemical  industry  out  of  a  change 
of  theoretical  interest  only.  The  formation  of  nitric  acid  from 
air  and  water  is  based  on  reactions  represented  by  the  following 
equations : 

N2  +  02  +  43,200  cals  *=*  2NO 

4NO  +  3O2  +  2H2O  =  4HNO3 

The  second  reaction  takes  place  readily;  if  nitric  oxide  and  an 
excess  of  air  are  passed  through  water  at  ordinary  temperatures 
nitric  acid  is  formed.  The  first  reaction  is  the  one  that  required 
extensive  study  before  it  could  be  utilized  for  large-scale  production. 
The  reaction  is  a  reversible  one,  and,  consequently,  to  discover 
the  most  favorable  conditions  for  its  use  the  equilibria  at  different 
temperatures  had  to  be  determined,  and  the  rate  of  the  reaction  at 
these  temperatures,  investigated.  Since  the  formation  of  nitric 
oxide  is  an  endothermic  reaction,  rise  in  temperature  causes  the 
equilibrium  to  shift  in  such  a  way  that  the  proportion  of  nitric 
oxide  in  the  resulting  gases  increases  (law  of  mobile  equilibrium). 
With  rise  in  temperature  the  rate  at  which  equilibrium  is  attained 
increases, 


324 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


Electrode 
for  Arc...,, 


The  study  of  the  equilibrium  from  these  two  points  of  view 
disclosed  the  best  practical  conditions  for  carrying  out  the 
reaction.  It  was  found,  for  example,  that  at  1538°  the  gases  in 
equilibrium  when  air  was  used  contained  0.37  per  cent  nitric  oxide, 
at  1922°,  0.97  per  cent,  and  at  2927°,  5.0  per  cent.  The  rate  at 
which  equilibrium  was  attained  at  the  high  temperatures  required 
was  measurable  but  high  enough  for  practical  purposes,  and  it 
dropped  off  rapidly  with  falling  temperature.  The  data  obtained 
led  to  the  conclusion  that  the  reaction  should  be  carried  out  at  the 
highest  temperature  possible,  that  no  catalytic  agent  was  neces- 
sary, and  that  after  the  gases  had  reached  equilibrium  they 
should  be  suddenly  cooled  in  order  to  prevent  the  shifting  of  the 
equilibrium,  which  would  result,  with  slowly  falling  temperature, 
in  the  decomposition  of  the  nitric  oxide  produced. 

A  number  of  forms  of  apparatus  have  been  devised  to  meet 
these  conditions.  Since  the  temperature  required  is  so  high,  an 

electric  arc  is  the  source  of 
heat  in  all  of  them.  In  the 
Birkeland-Eyde  process  used 
in  Norway,  an  arc  between 
two  carbon  poles  is  flattened 
out  into  a  fan-like  discharge 
by  means  of  an  electro-mag- 
net. A  cross-section  of  the 
apparatus  is  shown  in  Fig.  31. 
The  gases  from  the  furnace, 
after  cooling,  are  passed 
through  towers  filled  with  tiles 
down  which  water  trickles. 

Arrin/ef-   '^Sas  and  Air  Outlet        The  dilute  nitric  acid  obtained 
FlG-  31.  is  neutralized  with  lime,  CaO, 

and  is  converted  into  calcium 

nitrate,  which  is  obtained  on  evaporation  of  the  solution.     The 
yield  is  said  to  be  70  grams  of  nitric  acid  per  kilowatt  hour. 

It  is  of  interest  to  compare  here  the  process  employed  to  make 
nitrogen  and  hydrogen  unite  to  form  ammonia  (338)  with  that 
used  to  prepare  nitric  oxide  from  nitrogen  and  oxygen.  In  the 
first  case  the  reaction  is  exothermic  and  in  the  second  case  it  is 
endothermic;  and  this  difference  led  to  quite  different  procedures 


NITRIC  ACID,  NITROUS  ACID,  OXIDES  OF  NITROGEN  325 

in  using  the  reactions  for  the  two  syntheses.  The  nitrogen- 
hydrogen  reaction  was  carried  out  at  as  low  a  temperature  as  pos- 
sible, and  as  its  rate  decreased  as  the  temperature  was  lowered  it 
was  necessary  to  use  a  catalyst.  The  nitrogen-oxygen  reaction  was 
carried  out  at  as  high  a  temperature  as  possible,  and  as  increase 
in  temperature  is  associated  with  increase  in  rate,  no  catalyst  was 
required.  In  the  nitrogen-hydrogen  reaction  there  is  a  change  in 
volume  as  the  result  of  the  union — 4  volumes  become  2 — and,  as  a 
consequence,  the  synthesis  of  ammonia  was  carried  out  at  the 
highest  practical  pressure.  In  the  nitrogen-oxygen  reaction  there 
is  no  change  in  volume  and  the  synthesis  was  carried  out  at 
ordinary  pressures.  In  the  synthesis  of  nitric  oxide  energy  is 
absorbed  and  the  cost  of  electrical  power  is  a  factor  in  the  cost  of 
the  nitric  acid  manufactured.  In  the  synthesis  of  ammonia  energy 
is  given  off;  the  power  necessary  is  only  that  required  to  compress 
the  gases,  and  we  have  seen  that  a  part  of  this  is  recovered.  It  is 
important  to  emphasize  the  fact  that  it  was  the  application  of  the 
principles  underlying  mobile  equilibrium  that  made  it  possible 
to  develop  into  commercial  processes  reactions  which  under 
ordinary  circumstances  furnished  but  an  exceedingly  small  frac- 
tion of  a  per  cent  of  the  desired  product. 

359.  The  most  recently  developed  method  of  manufacturing 
nitric  acid  is  based  on  the  oxidation  of  ammonia.  The  economic 
significance  of  this  process  has  already  been  mentioned  and  it  was 
pointed  out  that  it  was  used  during  the  recent  war.  The  process 
is  simple  in  principle.  When  a  mixture  of  ammonia  and  air  is 
passed  over  a  proper  catalyzer,  nitric  oxide  is  formed  as  the  result 
of  a  reaction  between  the  ammonia  and  the  oxygen  in  the  air: 

4NH3  +  5O2  =  4NO  +  6H2O 

When  the  nitric  oxide  is  cooled  it  reacts  with  more  oxygen  and 
nitrogen  dioxide,  NO2,  is  formed;  the  latter,  when  mixed  with  air 
and  passed  into  water,  is  converted  into  nitric  acid : 

4NO2  +  2H2O  +  O2  =  4HNO3 

The  oxidation  of  ammonia  is  carried  out  in  chambers  lined  with 
fire-brick  in  which  several  layers  of  the  catalyzer  are  placed.  In 
some  plants  fine  platinum  gauze  is  used  as  the  contact  agent  and 
in  others  an  oxide  of  iron.  The  reaction  goes  to  completion  at  the 


326  INORGANIC  CHEMISTRY  FOR  COLLEGES 

temperature  used,  about  600°,  and  being  exothermic  does  not 
require  the  application  of  external  heat  after  it  has  been  started. 
The  nitric  oxide  is  cooled,  mixed  with  more  air,  and  passed  through 
towers  lined  with  tile  down  which  water  trickles.  In  order  to 
effect  complete  absorption  of  the  gases  several  towers  are  used,  the 
last  containing  a  solution  of  sodium  carbonate.  Several  hundred 
thousand  tons  of  nitric  acid  and  nitrates  have  been  made  in  a 
single  year  by  this  process.  . ;  ' 

360.  Physical  Properties  of  Nitric  Acid. — Nitric  acid  is  a  color- 
less liquid  which  boils  at  86°,  and  has  the  specific  gravity  1.56  at 
0°.     It  usually  has  a  yellow  color  which  is  produced  by  the  pres- 
ence of  oxides  of  nitrogen  formed  as  the  result  of  the  decomposition 
of  some  of  the  acid.     The  concentrated  nitric  acid  of  commerce  is 
the  constant  boiling  mixture  of  the  acid  and  water;    it  contains 
68.6  per  cent  nitric  acid,  has  the  specific  gravity  1.41  at  15°  and 
boils  at  120.5°. 

361.  Chemical  Behavior  of  Nitric  Acid. — The  acid  decomposes 
slowly  at  its  boiling-point  according  to  the  following  equation: 

4HNO3  =  2H2O  -f  4NO2  +  O2 

Nitric  acid  is  miscible  with  water  in  all  proportions  and  the 
solutions  show  strong  acidic  properties.  The  acid  is  dissociated 
into  H+  and  NO3~  ions,  and,  as  is  usual,  the  extent  to  which  the 
dissociation  takes  place  is  determined  by  the  concentration  of  the 
acid;  in  one-tenth  normal  solution  at  18°  it  is  92  per  cent  disso- 
ciated; the  ionization  of  hydrochloric  acid  at  this  dilution  is  92 
and  of  sulphuric  acid  61  per  cent.  The  ionization  in  a  solution 
containing  62  per  cent  of  nitric  acid  is  only  9  per  cent.  The  extent 
to  which  ionization  has  taken  place  in  a  solution  of  nitric  acid  is 
an  important  factor  in  its  chemical  behavior;  undissociated  nitric 
acid  is  an  active  oxidizing  agent,  whereas  the  ions  produced  from  it, 
NOa~,  do  not  show  this  property.  Anhydrous  nitric  acid  is  such 
a  powerful  oxidizing  agent  that  a  flame  is  sometimes  produced 
when  it  is  brought  into  contact  with  oxidizable  substances.  This 
can  be  readily  shown  by  simple  experiments.  If  warm  nitric 
acid  is  poured  on  sawdust  which  has  been  slightly  heated,  the  oxi- 
dation is  so  violent  that  the  wood  ignites  and  burns  with  a  flame. 
The  dense  brown  fumes  produced  are  nitrogen  dioxide  formed  from 


NITRIC  ACID,  NITROUS  ACID,  OXIDES  OF  NITROGEN  327 

the  nitric  acid.  If  a  bit  of  wool  is  placed  in  the  mouth  of  a  test- 
tube  containing  some  pure  nitric  acid,  and  the  latter  is  then  heated, 
the  vapors  given  off  from  the  acid  ignite  the  wool,  which  burns 
with  a  flame.  It  is  evident  that  great  care  must  be  used  in  ship- 
ping and  storing  pure  nitric  acid.  The  danger  of  fire  from  con- 
centrated nitric  acid  is  not  so  great,  but  care  is  necessary  with  it, 
nevertheless,  because  it  is  highly  corrosive. 

362.  The  reduction-product  formed  from  nitric  acid  when  it 
oxidizes  is,  as  we  would  conclude,  dependent  on  the  strength  of  the 
acid  and  the  activity  of  the  reducing  agent  with  which  it  interacts. 
In  most  cases,  however,  the  chief  reduction-product  is  nitric  oxide, 
the  acid  breaking  down  as  indicated  by  the  following  equation: 

2HNO3  =  2NO  +  H2O  +  3O 

It  must  be  emphasized  that  this  is  not  an  equation  expressing  the 
decomposition  of  nitric  acid  alone;  it  represents  the  decomposition 
which  takes  place  when  nitric  acid  is  brought  into  contact  with 
some  oxidizable  substance,  and  it  serves,  therefore,  as  a  partial 
equation  in  a  series  representing  the  oxidation  of  a  substance  by 
nitric  acid.  We  shall  see  that  the  equation  is  used  repeatedly 
when  oxidation  reactions  are  written. 

363.  Nitric  acid  oxidizes  all  the  metals  except  the  so-called 
noble  metals,    gold  and  platinum,    for  example.      The  nitrates 
formed  are  soluble  in  water,  and,  consequently,  nitric  acid  dis- 
solves these  metals.     It  was  on  account  of  this  solvent  action  of 
the  acid  that  it  was  called  by  the  alchemists  "  aqua  fortis."     In 
writing  equations  for  the  reactions  involved,  it  is  best  to  separate 
them  into  steps  and  combine  the  partial  equations  in  the  way 
already  explained  (287,  288).     For  example,  the  oxidation  of  cop- 
per by  nitric  acid  can  be  represented  as  follows ; 

2HNO3  =  2NO  +  H2O  +  [3O] 
3Cu  +  [30]  =  [3CuO] 

[3CuO]  +  6HNO3  =  3Cu(N03)2  +  3H20 
3Cu  +  8HNO3  =  3Cu(NO3)2  +.4H2O  +  2NO 


328  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  oxidation  of  silver,  which  has  the  valence  1,  can  be  written  in 
a  similar  way: 

2HNO3  =  2NO  +  H2O  +  [3O] 

6Ag  +  [30]  =  [3Ag20] 

[3Ag2O]  +  6HNO3  =  6AgNO3  +  3H2O 

6Ag  +  8HNO3  =  6AgN03  +  4H2O  +  2NO 

In  this  case  all  the  integers  indicating  the  number  of  molecules  can 
be  divided  by  two,  and  the  equation  is  simplified  to  read  as  follows : 

3Ag  +  4HNO3  =  3AgNO3  +  2H2O  +  NO 

Many  acid-forming  elements  are  oxidized  by  strong  nitric  acid; 
the  equation  for  the  reaction  with  sulphur  is  as  follows : 

2HN03  =  2NO  +  [H20]  +  [3O] 

S  +  [30]  =  [S03] 

[S03]  +  [H20]  =  H2S04 

S  +  2HNO3  =  2NO  +  H2SO4 

Carbon  when  heated  with  strong  nitric  acid  is  oxidized: 
4HNO3  =  4NO  +  2H2O  +  [6O] 

3C  +  [6O]  =  3CO2 
3C  +  4HNO3  =  4NO  +  2H2O  +  3CO2 

The  nitric  oxide  formed  in  all  the  reactions  mentioned  is  converted 
into  nitrogen  dioxide  if  air  is  present,  because  the  colorless  nitric 
oxide  unites  at  ordinary  temperatures  with  the  oxygen  of  the  air 
to  form  nitrogen  dioxide,  which  is  a  brown  gas.  If  the  oxidations 
are  carried  out  in  the  presence  of  an  excess  of  concentrated  nitric 
acid,  the  dioxide  is  also  formed. 

When  the  most  active  metals  react  with  nitric  acid,  hydrogen 
is  formed;  if  the  solution  is  very  dilute  and  the  acid  almost  com- 
pletely in  the  form  of  ions,  the  hydrogen  escapes;  in  slightly 
stronger  solutions  the  hydrogen  is  oxidized  by  the  nitric  acid  to 
water  and  the  acid  is  reduced  to  ammonia,  which  forms  ammonium 
nitrate  with  some  of  the  acid  present.  Under  special  conditions 
nitric  acid  when  it  acts  as  an  oxidizing  agent  may  be  reduced  to 


NITRIC  ACID,  NITROUS  ACID,  OXIDES  OF  NITROGEN    329 

nitrous  oxide,  hydroxylamine,  or  hydrazine  (352) ;  the  reaction  to 
be  definitely  remembered,  however,  is  its  reduction  to  nitric  oxide. 
Many  compounds  are  oxidized  by  the  acid;  the  reaction  in  the 
case  of  sulphur  dioxide  is  expressed  by  the  following  equations : 

2HNO3  =  H2O  +  2NO  +  [3O] 

3S02  +  [30]  =  [3S03] 
[3S03]  +  3H20  =  3H2S04 


2HN03  +  3SO2  +  2H2O  =  2NO  +  3H2SO4 

When  nitric  acid  oxidizes  hydrochloric  acid,  the  nitric  oxide 
formed  unites  with  some  of  the  chlorine  produced  by  the  oxidation 
to  form  nitrosyl  chloride,  NOC1: 

2HN03  =  H2O  +  [2NO]  +  [3O] 
6HC1  +  [3O]  =  3H2O  +  6C1 
[2NO]  +  2C1  -  2NOC1 


2HNO3  +  6HC1  -  4H2O  +  2C12  +  2NOC1 
HNO3  +  3HC1  =  2H2O  +  C12  +  NOC1 

It  will  be  seen  from  the  equations  that  chlorine  is  formed.  Since 
chlorine  reacts  with  many  substances  and  converts  them  into 
chlorides  which  are  soluble,  the  mixture  of  hydrochloric  acid  and 
nitric  acid  is  an  excellent  solvent.  The  mixture  was  called  aqua 
regia  by  the  alchemists  because  it  dissolved  the  royal  metal,  gold. 
Aqua  regia  is  much  used  in  the  chemical  laboratory  to  convert 
various  insoluble  substances  into  soluble  compounds.  It  is  gen- 
erally prepared  by  mixing  1  volume  of  concentrated  nitric  acid  and 
3  volumes  of  concentrated  hydrochloric  acid. 

364.  A  solution  of  nitric  acid  in  water  shows  all  the  charac- 
teristic properties  of  an  acid;  it  reacts  with  oxides  and  hydroxides 
of  metals  to  form  nitrates.  There  are  certain  organic  compounds 
called  alcohols  which  contain  the  hydroxyl  group  and  interact 
with  strong  acids  to  form  salts.  Common  grain  alcohol  has  the 
formula  C2H5OH — it  is  the  hydroxide  of  the  radical  C2H5,  which 
is  called  ethyl.  Alcohol  reacts  with  nitric  acid  and  forms  ethyl 
nitrate : . 

C2H5OH  +  HN03  =  C2H5N03  +  H2O 


330  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Since  alcohol  can  be  oxidized  by  nitric  acid,  special  precautions  are 
necessary  in  carrying  out  the  reaction  to  prevent  the  oxidation 
which  often  takes  place  with  explosive  violence.  Glycerine  is  an 
alcohol  of  the  formula  C3H5(OH)3  and  when  treated  with  a  mixture 
of  nitric  and  sulphuric  acids  forms  a  nitrate,  C3H5(NO3)3,  which 
is  commonly  called  nitroglycerine.  The  product,  which  is  a  liquid, 
is  made  in  large  quantities  and,  mixed  with  sawdust,  clay,  and 
other  substances  is  used  under  the  name  of  dynamite  as  an  ex- 
plosive. 

Cellulose,  which  is  the  chief  constituent  of  raw  cotton  and 
wood,  is  also  an  alcohol.  It  forms  a  nitrate  with  nitric  acid : 

C6H702(OH)3  +  3HNO3  =  C6H7O2(NO3)3  +  3H2O 

The  nitrate  is  commonly  called  nitrocellulose  and  is  the  chief 
ingredient  of  smokeless  powder. 

Nitric  acid  reacts  with  many  hydrocarbons  and  their  deriva- 
tives and  forms  the  so-called  nitro  compounds,  all  of  which  contain 
the  characteristic  nitro  group,  N02.  For  example,  it  reacts  with 
benzene  as  follows: 

C6H6  +  HNO3  =  C6H5NO2  +  H2O 

The  product,  nitrobenzene,  can  be  reduced  by  nascent  hydrogen 
to  aniline,  CeHsNH^,  which  is  used  in  the  manufacture  of  the 
so-called  aniline  dyes. 

With  toluene,  a  derivative  which  contains  three  nitro  groups 
can  be  made: 

CH3.C6H5  +  3HN03  -  CH3.C6H2(N02)3  +  3H2O 

Trinitrotoluene,  commonly  called  T.N.T.,  is  a  very  important 
high  explosive  used  in  shells  for  military  purposes.  Picric  acid, 
which  is  a  valuable  explosive,  is  also  used  in  large  quantities  in 
warfare.  It  is  trinitrophenol,  C6H2(OH)(NO2)3,  and  is  made 
from  phenol,  CeBkOH,  which  is  often  called  carbolic  acid. 

365.  When  all  the  explosives  which  have  been  mentioned  de- 
compose, there  is  a  liberation  of  the  energy  which  was  stored  up  in 
the  compound  as  the  result  of  the  union  of  nitrogen  and  oxygen  in 
the  nitric  acid  from  which  they  were  prepared ;  the  carbon  present  is 
oxidized  to  carbon  dioxide  and  carbon  monoxide,  or  if  there  is  not 
enough  oxygen  present  in  the  molecule  to  oxidize  all  the  carbon, 


NITRIC  ACID,  NITROUS  ACID,  OXIDES  OF  NITROGEN  331 

it  is  liberated  as  such  and  appears  as  black  smoke.  Some  explo- 
sives, like  nitroglycerine,  decompose  very  readily;  others  can  be 
handled  with  safety,  since  it  is  necessary  to  detonate  them  before 
they  explode.  A  detonater  is  a  compound  that  explodes  when  it  is 
subjected  to  a  shock  such  as  the  blow  of  the  hammer  of  a  rifle. 
The  decomposition  which  takes  place  sets  up  vibrations  which 
cause  the  explosion  of  compounds  not  readily  influenced  by  simple 
mechanical  shocks. 

366.  Hair,  flesh,  feathers,  silk,  and  other  nitrogenous  constit- 
uents of  living  things,  which  are  called  proteins,  are  formed  as  the 
result  of  the  union  of  a  large  number  of  complex  molecules  one  of 
which  is  derived  from  benzene.     When  any  of  these  substances  is 
treated  with  strong  nitric  acid,  a  nitro  compound  is  formed  just  as 
benzene  yields   nitrobenzene.     The   compound,   like  most  nitro 
compounds,  is  yellow.     It  is  for  this  reason  that  nitric  acid  pro- 
duces a  yellow  stain  when  left  in  contact  with  the  skin  or  woolen 
fabrics.     When  acids  are  allowed  to  stay  in  contact  with  dyed 
cloth  a  red  stain  generally  develops  as  a  result  of  a  chemical  change 
in  the  dye.     Such  a  stain  can  be  removed  by  treating  it  with 
ammonia;  the  dye  in  this  case  acts  as  an  indicator  and  the  effect  of 
an  acid  is  neutralized  by  the  application  of  a  base.     Ammonia  is 
used  for  this  purpose  because  it  does  not  materially  affect  the 
fabric,  and  it  is  volatile  and  any  excess  is  removed  by  evaporation. 
Sodium  hydroxide  decomposes  proteins  and,  therefore,  attacks 
wool  and  silk.     When  a  stain  produced  by  nitric  acid  is  treated 
with  ammonia  the  original  color  is  not  restored  because  the  yellow 
nitro  compound  produced  by  the  acid  is  not  destroyed. 

367.  Nitrates. — All  normal  nitrates  are  soluble  in  water,  and 
the  difficultly  soluble  basic  nitrates  are  soluble  in  the  presence  of 
an  excess  of  nitric  acid.     All  nitrates  decompose  when  heated, 
those  of  the  more  active  elements  requiring  the  higher  temperatures 
to  effect  decomposition.     In  the  case  of  sodium  and  potassium, 
the  nitrates  can  be  melted  without  decomposition,  but  at  red 
heat  they  lose  oxygen  and  pass  into  nitrites;   sodium  nitrite  may 
be  formed  in  this  way: 

2NaNO3  =  2NaNO2  +  02 

The  nitrates  of  all  metals  except  those  that  resemble  sodium  in 
chemical  properties  break  down  into  the  oxide  of  the  metal  and 


332  INORGANIC  CHEMISTRY  FOR  COLLEGES 

the  anhydride  of  the  acid.     The  anhydride  of  nitric  acid  is  NoOs: 
2HNO3  =  N2O5  +  H2O 

Nitric  acid  anhydride  is  unstable  at  the  temperature  at  which 
nitrates  decompose  and  breaks  down  into  nitrogen  dioxide  and 
oxygen: 

2N2O5  -  4NO2  +  O2 

The  equation  which  represents  the  decomposition  of  copper  nitrate 
when  heated  is,  therefore: 

2Cu(NO3)2  =  2CuO  +  4NO2  +  O2 

368.  The  test  for  nitrates  is  based  on  the  production  of  a  color, 
and  not,  as  usual,  on  the  formation  of  some  insoluble  compound,  as 
is  the  case  with  sulphates,  chlorides,  etc.  The  formation  of  a 
characteristic  insoluble  compound  cannot  be  used  in  this  case 
because  of  the  fact  that  nitrates  are  soluble.  We  have  learned 
that  when  nitric  acid  oxidizes  substances  it  is  reduced  to  nitric 
oxide.  This  oxide  forms  a  compound  with  ferrous  sulphate, 
FeSO4,  which  is  highly  colored  in  great  dilution,  and  its  formation, 
therefore,  serves  as  a  delicate  test  for  nitric  oxide  and  for  nitric 
acid  provided  the  latter  is  brought  into  contact  with  something 
that  reduces  it  to  nitric  oxide.  Ferrous  sulphate  is  oxidized  by 
nitric  acid  to  ferric  sulphate,  so  it  serves  two  purposes  in  the 
test  based  on  these  reactions.  The  test  is  carried  out  as  fol- 
lows: A  solution  of  the  substance  to  be  tested  is  mixed  in  a 
test-tube  with  a  strong  solution  of  ferrous  sulphate.  The  tube 
is  inclined  and  concentrated  sulphuric  acid  cautiously  poured  in. 
As  the  acid  flows  down  the  side  of  the  tube  it  sinks  to  the 
bottom  and  forms  a  heavy  layer  beneath  the  aqueous  solution. 
If  a  nitrate  is  present  a  brown  ring  appears  at  the  juncture 
of  the  two  liquids.  The  sulphuric  acid  liberates  nitric  acid  from 
the  nitrate;  the  acid  oxidizes  a  part  of  the  ferrous  sulphate  and  is 
thereby  reduced  to  nitric  oxide,  and  finally  the  nitric  oxide  unites 
with  some  of  the  ferrous  sulphate  to  form  a  compound  of  the 
formula  FeSCU  •  NO,  which  possesses  a  brown  color.  The  test  is  a 
delicate  one  because  it  shows  the  presence  of  a  very  small  amount 
of  nitric  acid  or  a  nitrate. 

The  individual  salts  of  nitric  acid  will  be  described  in  connec- 


NITRIC  ACID,  NITROUS  ACID,  OXIDES  OF  NITROGEN  333 

tion  with  the  discussion  of  the  chemistry  of  the  metals  which  they 
contain.  Ammonium  nitrate  was  used  in  large  quantities  as  a 
high  explosive  during  the  recent  war.  The  compound  under 
ordinary  circumstances  is  inert.  When  the  dry  salt  is  heated  it 
decomposes  into  nitrous  oxide  and  water, 

NH4NO3  =  N2O  +  2H2O, 

and  the  reaction  serves  as  the  most  convenient  method  of  pre- 
paring the  gas.  When,  however,  the  salt  is  detonated  it  decom- 
poses with  explosive  violence.  One  of  the  most  valuable  fillings 
for  high  explosive  shells  is  a  mixture  of  ammonium  nitrate  and 
T.N.T.;  it  is  called  amatol. 

NITROUS  ACID 

369.  Nitrous  acid  is  used  in  making  the  so-called  azo  dyes,  and 
as  these  are  readily  prepared  and  furnish  a  great  variety  of  colors 
and  shades,  large  amounts  of  the  salts  of  the  acid  are  manufactured. 
The  acid  itself  is  very  unstable,  as  it  breaks  down  largely  into  its 
anhydride   and   water   when   liberated.     The    chemical   relation 
between  nitrous  acid,  HNC>2,  and  nitric  acid,  HNOs,  is  similar 
to  that  between  sulphurous  acid,   H^SOs,   and  sulphuric  acid, 
H2SO4,  the  similarity  being  traceable  to  the  fact  that  in  the  two 
cases  the  acids  differ  in  composition  by  one  oxygen  atom.    We  shall 
see  later  that  there  are  many  cases  in  which  one  element  forms  two 
acids  containing  different  amounts  of  oxygen,  and  that  the  differ- 
ence in  their  chemical  properties  can  be  traced  to  this  cause.     It 
is  well  to  emphasize  the  relationship  here. 

370.  The  methods  commonly  used  to  prepare  nitrous  acid  and 
nitrites  are  not  analogous  to  those  used  in  the  case  of  the  sulphur 
compounds  for  a  number  of  reasons.    The  supply  of  raw  materials 
is  different  in  the  two  cases,  and  as  the  reactions  are  endothermic 
in  one  case  and  exothermic  in  the  other,  the  ease  of  preparation  of 
the  nitrogen  compounds  and  the  sulphur  compounds  from  the 
elements  is  markedly  different.     It  will  be  recalled  that  sulphur 
is  the  raw  material  from  which  sulphurous  acid  and  sulphuric 
acid  are  made,  and  that  it  is  only  recently  that  nitrogen  has  been 
used  as  the  source  of  nitric  acid.     Nitrous  acid  and  its  salts  are 
now  manufactured  from  the  oxides  produced  by  burning  nitrogen, 


334  INORGANIC  CHEMISTRY  FOR  COLLEGES 

but  the  older  method  starting  with  the  sodium  nitrate  that  occurs 
ready  made  in  nature  is  the  one  commonly  used. 

When  sodium  or  potassium  nitrate  is  heated  to  a  high  tempera- 
ture a  part  of  the  oxygen  is  lost,  and  a  nitrite  is  formed : 

2NaNO3  =  2NaNO2  +  O2 

In  manufacturing  nitrites  lead  is  usually  stirred  with  the  molten 
mass;  it  helps  in  the  reduction  by  uniting  with  the  oxygen,  and 
is  converted  into  lead  oxide,  PbO.  When  the  reaction  is  complete 
the  mixture  is  allowed  to  cool  and  is  treated  with  hot  water;  the 
solution  on  evaporation  yields  sodium  nitrite  in  the  form  of 
colorless  crystals. 

371.  When  a  solution  of  sodium  nitrite  is  treated  with  an  acid  a 
reaction  takes  place  which  is  analogous  to  that  in  the  case  of 
sodium  sulphite,  and  the  reasons  underlying  the  double  decompo- 
sition are  the  same  in  the  two  cases  (287) .  Nitrous  acid  is  formed 
and  in  part  decomposes  into  nitrous  anhydride  and  water; 

2NaNO2  +  H2SO4  =  Na2SO4  +  2HNO2 
2HNO2  <=>  N2O3  +  H2O 

The  nitrous  anhydride  escapes  as  a  brown  gas. 

When  nitrous  anhydride  is  passed  into  a  solution  of  sodium 
hydroxide  it  reacts  and  forms  a  nitrite,  just  as  sulphur  dioxide 
under  the  same  conditions  forms  a  sulphite : 

H2O  +  N2O3  <=*  2HNO2 
NaOH  +  HNO2  =  NaNO2  +  H2O 

Like  sulphurous  acid,  nitrous  acid  is  a  mild  oxidizing  and  reducing 
agent.  The  reaction  which  takes  place  between  hydriodic  acid 
and  sulphurous  acid  (sulphur  dioxide  and  water)  has  already  been 
given : 

4HI  +  SO2  =  2I2  +  2H2O  +  S 

The  reaction  between  hydriodic  acid  and  nitrous  acid  (nitrous 
anhydride  and  water)  is  similar  to  the  one  just  given,  although 
nitrogen  is  not  formed  in  this  case: 

2HI  +  N203  =  I2  +  H20  +  2NO 


NITRIC  ACID,  NITROUS  ACID,  OXIDES  OF  NITROGEN    335 

The  oxidation  of  hydriodic  acid  can  be  traced  to  the  same  cause 
in  the  two  cases.  Both  nitrogen  and  sulphur  can  exist  in  com- 
pounds in  which  the  elements  have  lower  valencies  than  in 
sulphurous  and  nitrous  anhydrides.  In  sulphur  dioxide  and  sul- 
phurous acid  sulphur  has  the  valence  4,  and  when  these  com- 
pounds serve  as  oxidizing  agents  the  valence  of  the  sulphur  falls 
to  zero  when  free  sulphur  is  formed.  When  nitrous  acid  or  its 
anhydride  act  as  oxidizing  agents  they  do  not  give  up  all  the 
oxygen  with  which  the  nitrogen  is  combined,  but  are  reduced  to 
nitric  oxide,  the  change  in  valence  being  from  3  to  2. 

372.  When  sulphur  dioxide  acts  as  a  reducing  agent  it  unites 
with  more  oxygen  and  sulphur  trioxide  (sulphuric  acid)  is  formed, 
the  valence  change  being  from  4  to  6 :  SO2  — >  SOs.     When  nitrous 
anhydride  serves  as  a  reducing  agent  it  passes  to  nitric  anhydride 
(nitric  acid)  and  the  valence  change  is  from  3  to  5:  N2Oa  — »  ^Os. 
Sulphurous  and  nitrous  acids  act  as  reducing  agents  only  when 
brought  into  contact  with  powerful  oxidizing  agents.     When  an 
acidified  solution  of  potassium  permanganate  is  treated  with  either 
substance  the  purple  color  of  the  salt  disappears;    the  perman- 
ganate is  reduced  to  a  colorless  compound  and  the  sulphurous 
acid  is  oxidized  to  sulphuric  acid  or  the  nitrous  acid  to  nitric  acid. 

While  there  is  a  striking  analogy  between  the  nitrogen  and 
sulphur  compounds,  there  is  a  difference  between  them  in  their 
activity  as  reducing  or  oxidizing  agents;  sulphurous  anhydride  is  a 
more  active  reducing  agent  than  nitrous  anhydride,  and  nitric 
acid  a  more  active  oxidizing  agent  than  sulphuric  acid. 

373.  Nitrites. — The  nitrites  of  most  of  the  metals  are  known; 
they  are  all  soluble  in  water,  and  yield  nitrous  anhydride  when 
treated  with  an  acid.     The  formation  of  this  gas  is  used  in  the 
test  for  nitrites.     The  solution  is  acidified  and  warmed  if  necessary; 
if  a  nitrite  is  present'  a  yellow-brown  gas  is  evolved.     When  the 
amount  of  nitrite  present  is  small  and  the  result  indefinite  on 
account  of  the  small  amount  of  gas  given  off,  the  liquid  or  the  vapor 
from  it  is  tested  with  a  bit  of  paper  which  has  been  moistened  with 
a  solution  containing  potassium  iodide  and  starch.     The  nitrous 
anhydride  oxidizes  the  hydriodic  acid  and  sets  free  iodine  from  the 
iodide,  according  to  the  reaction  which  has  already  been  explained 
(371) ,  and  the  liberated  iodine  produces  with  the  starch  a  charac- 
teristic blue  color.     The  presence  of  the  starch  makes  the  test 


336  INORGANIC  CHEMISTRY  FOR  COLLEGES 

more  delicate,  for  an  amount  of  iodine  that  could  not  be  recognized 
by  its  color  produces  a  marked  blue  with  starch.  Nitrites,  as 
might  be  expected,  are  much  more  stable  under  the  influence  of 
heat  than  nitrates.  The  nitrites  of  potassium  and  sodium  resist 
very  high  temperatures;  those  of  the  heavy  metals  decompose 
at  red  heat  or  lower  and  yield  oxides  of  the  metals  and  nitric 
oxide  and  oxygen,  which  are  produced  as  the  result  of  the  de- 
composition of  nitrous  anhydride. 

374.  Water  Analysis. — Having  acquired  a  knowledge  of  the 
chemistry  of  ammonia,  nitrous  acid,  and  nitric  acid  we  are  now  in  a 
position  to  understand  the  significance  of  the  methods  used  in  the 
chemical  analysis  of  water  for  sanitary  purposes.  One  of  the  chief 
sources  of  contamination  of  water  is  sewage,  which  consists  largely 
of  animal  refuse  matter.  A  study  of  the  amount  of  this  present 
in  a  sample  of  water  and  the  changes  it  has  undergone  gives  a  good 
indication  of  the  purity  of  the  water  and  its  value  for  household 
use.  The  complex  organic  nitrogen  compounds  undergo  decom- 
position slowly  as  the  result  of  oxidation  which  is  brought  about 
through  the  influence  of  certain  bacteria  always  present.  This 
process  of  nitrification  takes  place  in  steps;  first  ammonia  is  set 
free  by  one  kind  of  nitrifying  bacteria,  then  this  is  oxidized  to 
nitrous  acid  by  another  kind,  and,  finally,  the  nitrous  acid  is  con- 
verted into  nitric  acid  by  a  third  kind.  As  these  processes  occur 
slowly  it  is  possible,  by  determining  quantitatively  the  amount  of 
undecomposed  material,  ammonia,  nitrous  acid,  and  nitric  acid 
present  in  a  sample  of  water,  to  estimate  the  extent  to  which  con- 
tamination has  taken  place,  and  whether  it  has  been  recent  or  not. 
By  the  time  all  the  nitrogen  has  been  converted  into  nitrates  the 
impurities  have  disappeared  and  the  water  may  be  considered  safe. 

In  analyzing  a  sample  of  water  a  measured  quantity  of  it  is 
first  distilled,  and  the  amount  of  the  free  ammonia  which  passes 
over  with  the  steam  is  determined.  Potassium  permanganate, 
KMnO4,  is  next  added  and  the  distillation  continued.  Potassium 
permanganate  oxidizes  the  organic  matter,  and  the  nitrogen  present 
is  liberated  as  ammonia,  the  amount  of  which  is  determined  as 
before.  This  ammonia  is  a  measure  of  the  undecomposed  organic 
matter  present  in  the  water  and  is  called  albuminoid  ammonia,  on 
account  of  the  fact  that  it  is  derived  from  the  so-called  albumens 
present.  Nitrites  and  nitrates  are  determined  in  separate  samples  of 


NITRIC  ACID,  NITROUS  ACID,  OXIDES  OF  NITROGEN     337 

the  water  under  examination.  All  the  quantitative  determinations 
are  made  with  the  aid  of  substances  that  produce  colors  when 
brought  into  contact  with  the  compound  tested  for.  Colorimetric 
methods  are  necessary  in  the  analysis  because  such  small  quanti- 
ties of  the  materials  to  be  determined  are  present;  we  can  see  a 
color  distinctly  when  it  would  be  impossible  to  weigh  the  minute 
amount  of  material  producing  the  color.  The  amount  of  the 
nitrogen  compound  present  is  determined  by  the  depth  of  color 
produced.  In  each  case  the  solution  is  placed  in  a  tube  and  the 
intensity  of  the  color  compared  with  a  series  of  standards  made  up 
from  samples  of  water  to  which  have  been  added  known  quan- 
tities of  nitrites  or  nitrates.  Below  are  given  the  results  of  the 
analysis  of  a  sample  of  the  water  supplied  to  New  York  City. 

ANALYSIS  OF  CROTON  WATER 

Parts 
per  million 

Free  Ammonia 0 . 015 

Albuminoid  ammonia 0 . 170 

Nitrogen  in  nitrites 0 . 000 

Nitrogen  in  nitrates 0 . 250 

Chlorine  in  chlorides 2 . 100 

Mineral  matter 66.000 

Total  solids 81 .000 

The  chlorine  which  is  present  in  a  water  chiefly  as  sodium  chloride 
is  determined,  as  it  is  also  an  indication  of  contamination  if  exces- 
sive amounts  are  found. 

OXIDES  OF  NITROGEN 

375.  There  are  five  oxides  of  nitrogen  which  form  a  complete 
series  in  which  nitrogen  shows  the  valence  of  1  to  5  inclusive; 
their  formulas  are  N2O,  NO,  N2O3,  N02,  and  N2O5.     In  addition, 
there  is  an  oxide  of  the  formula  N2O4  which  is  a  polymer  of  nitrogen 
dioxide,  NO2. 

376.  Nitrous  Oxide. — All  of  the  oxides  of  nitrogen   can  be 
prepared  from  nitric  acid.     With  the  exception  of  nitrogen  pen- 
toxide,  which  is  obtained  by  the  elimination  of  water  from  the  acid, 
the  compounds  are  made  by  reduction.     By  varying  the  concen- 
tration of  the  acid  by  dilution  with  water,  and,  therefore,  its  oxi- 
dizing power,  and  by  selecting  reducing  agents  of  increasing  activity 


338  INORGANIC  CHEMISTRY  FOR  COLLEGES 

we  can  prepare  the  oxides  in  which  nitrogen  has  the  valence  4,  3, 
2,  or  1.  This  method  is  not  the  most  convenient  one  in  all  cases, 
however.  Nitrous  oxide  is  formed  along  with  other  oxides  when 
dilute  nitric  acid  is  treated  with  zinc.  Priestley,  who  discovered 
the  gas  in  1772,  made  it  by  reducing  nitric  oxide  with  iron  filings 
in  the  presence  of  water.  The  most  convenient  method  of  prepara- 
tion is  to  heat  ammonium  nitrate  cautiously,  and  to  collect  the 
gas  over  water: 

NH4NO3  =  N2O  +  2H2O 

Nitrous  oxide  is  a  colorless  gas;  its  solubility  at  0°  is  130  vol- 
umes in  100  volumes  of  water,  and  at  25°,  60  in  100;  when  liquefied 
it  boils  at  -89.8°.  The  compressed  gas  is  furnished  in  cylinders 
and  is  used  as  an  anesthetic.  The  discovery  that  nitrous  oxide 
produces  unconsciousness,  or  when  breathed  in  smaller  quantities 
causes  hysterical  laughing,  was  made  in  1774  by  Humphrey  Davy. 
"  Laughing  gas  "  has  been  for  a  long  time  the  "  gas  "  adminis- 
tered by  dentists.  A  mixture  of  nitrous  oxide  and  oxygen  is 
generally  used,  since  by  varying  the  proportions  of  the  two  gases 
a  mixture  can  be  attained  which  produces  the  desired  degree  of 
insensibility. 

Nitrous  oxide  is  an  endothermic  compound;  when  it  decom- 
poses 18,000  calories  are  liberated  by  each  gram-molecule  of  the 
substance.  The  decomposition  can  be  brought  about  by  heat 
alone  or  better  by  a  detonator,  such  as  mercury  fulminate.  It  will 
be  recalled  that  ammonium  nitrate  is  a  powerful  explosive  when 
detonated;  its  use  for  this  purpose  is  based  upon  the  fact  that 
when  it  decomposes,  nitrogen  and  oxygen  are  set  free  with  the  evo- 
lution of  a  large  amount  of  heat.  Ammonium  nitrate  can  be 
broken  down  into  nitrous  oxide  and  water  by  gentle  heat  without 
explosion;  it  is  evident  that  the  explosive  property  of  the  salt  is 
due  to  the  breaking  apart  of  the  nitrogen  and  oxygen  in  the 
nitrous  oxide  formed  from  it. 

377.  Nitrous  oxide  decomposes  at  the  temperature  at  which 
most  substances  burn,  and  as  oxj^gen  is  liberated  the  gas  supports 
combustion.  When  nitrous  oxide  decomposes  into  its  elements  2 
volumes  of  the  gas  yield  2  volumes  of  nitrogen  and  1  volume  of 
oxygen: 

.      2N2O  =  2N2  +  O2 

Since  the  product  is  one-third  oxygen,  whereas  air  contains  but 


NITRIC  ACID,  NITROUS  ACID,  OXIDES  OF  NITROGEN  339 

one-fifth  oxygen,  it  is  clear  why  substances  burn,  in  general,  with  a 
more  brilliant  flame  in  nitrous  oxide  than  in  air.  The  proportion 
of  oxygen  is  higher  and  the  amount  of  inert  gas  to  be  heated  is 
consequently  less;  and  the  decomposition  of  the  nitrous  oxide 
produces  a  large  amount  of  heat.  These  factors  lead  to  the  pro- 
duction of  a  higher  temperature,  which  results  in  increased  lumi- 
nosity of  the  glowing  gases  in  the  flame. 

Nitrous  oxide  does  not  lose  its  oxygen  at  ordinary  temperatures 
and,  therefore,  it  does  not  affect  metals  as  oxygen  does;  iron  will 
not  rust  in  the  moist  oxide  as  it  does  in  air.  For  the  same  reason 
nitrous  oxide  does  not  convert  nitric  oxide  into  nitrogen  dioxide 
at  ordinary  temperatures.  It  will  be  recalled  that  oxygen  will 
bring  about  this  reaction.  We  could  readily  distinguish  nitrous 
oxide  from  air  or  oxygen  by  mixing  it  with  some  nitric  oxide;  if 
free  oxygen  is  present  the  colorless  nitric  oxide  is  converted  into 
the  dioxide,  which  has  a  yellow-brown  color. 

378.  We  see  from  the  behavior  of  nitrous  oxide  that  an  endo- 
thermic  compound  may  be  stable  although  it  contains  a  large 
amount  of  bound  up  chemical  energy.     In  order  to  make  this 
energy  available  it  is  necessary  to  bring  about  the  conditions 
which  result  in  the  disruption  of  the  molecules.     We  can  do  this 
by  raising  the  temperature  of  the  compound.     We  know  that  we 
heat  substances  to  make  them  interact,  and  it  is  reasonable  to 
believe  that  before  the  atoms  rearrange  themselves  to  form  new 
combinations  it  is  first  necessary  to  overcome  the  forces  holding 
the  atoms  together  in  the  molecule.     As  the  temperature  is  raised 
the    motion    of    the    atoms    in    the    molecules    increases    until 
finally  a  condition  is  reached  which  leads  to  the  separation  of  the 
atoms;  and  when  this  occurs  heat  is  given  off  or  absorbed,  as  the 
case  may  be.     The  temperature  at  which   decomposition  first 
takes    place    varies    with    different    substances.     The    study    of 
detonators  has  brought  out  the  fact  that  vibrations  set  up  by  one 
substance  undergoing  decomposition  may  induce  similar  vibrations 
within  the  molecules  of  another  substance  and  thus  cause  the 
latter  to  decompose  into  its  elements.     This  cannot  take  place,  of 
course,  when  the  decomposition  of  the  molecule  into  its  elements  is 
associated  with  the  absorption  of  energy. 

379.  Nitric  Oxide. — The  importance  of  nitric  oxide  in  chemical 
industry  has  been  repeatedly  emphasized;   it  serves  as  the  cata- 


340  INORGANIC  CHEMISTRY  FOR  COLLEGES 

lytic  agent  in  the  chamber  process  for  sulphuric  acid,  and  its  prepa- 
ration is  the  first  step  in  the  manufacture  of  nitric  acid  from  nitro- 
gen. The  gas  is  formed  when  nitric  acid  is  brought  into  contact 
with  substances  which  it  can  oxidize,  and  has  been  known  for  a  long 
time.  Priestley  tested  air  by  mixing  nitric  oxide  with  it;  it 
united  with  the  oxygen  present  and  formed  nitrogen  dioxide,  which 
could  be  dissolved  in  a  solution  of  sodium  hydroxide.  It  was, 
therefore,  possible  by  measuring  the  decrease  in  volume  of  the 
air  to  find  out  how  much  oxygen  it  contained. 

Nitric  oxide  is  most  readily  prepared  by  treating  copper  with 
concentrated  nitric  acid  which  has  been  diluted  with  an  equal  vol- 
ume of  water.  If  the  acid  is  not  diluted,  a  part  of  the  nitric  oxide 
formed  is  oxidized  by  it  to  nitrogen  dioxide.  The  equation  for  the 
reaction  has  already  been  discussed  in  detail  (363) ;  it  is  as  follows : 

3Cu  +  8HNO3  =  3Cu(NO3)2  +  2NO  +  4H29 

The  oxide  prepared  in  this  way  is  not  pure,  because  it  contains 
more  or  less  nitrous  oxide.  It  can  be  separated  from  the  latter  by 
passing  it  into  a  saturated  solution  of  ferrous  sulphate.  The 
nitric  oxide  forms  a  compound  with  the  sulphate  and  is,  there- 
fore, absorbed  by  the  solution.  The  brown  nitrosyl  ferrous  sul- 
phate is  the  substance  formed  in  the  test  for  nitric  acid  (368). 
When  the  solution  is  heated  nitric  oxide  is  evolved. 

380.  Nitric  oxide  is  a  colorless  gas,  which  is  liquefied  with  dif- 
ficulty and  is  very  slightly  soluble  in  water.  The  liquefied  gas 
boils  at  —153.6°.  The  thermochemistry  of  nitric  oxide  has  been 
discussed  in  some  detail  (358).  It  is  an  endothermic  compound 
which  contains  a  large  amount  of  bound-up  chemical  energy. 
A  comparison  of  the  thermochemical  equations  for  the  decompo- 
sition of  nitrous  oxide  and  nitric  oxide  is  of  interest: 

2N2O  =  2N2  +  O2  +  36,000  cals. 
2NO  =  N2  +  O2  +  43,200  cals. 

Like  nitrous  oxide,  nitric  oxide  can  be  detonated  and  an  explosion 
results.  It  is  more  stable  than  nitrous  oxide.  A  higher  tempera- 
ture is  required  to  decompose  it,  and,  as  a  consequence,  it  does  not 
support  combustion  so  readily.  If  a  burning  candle  is  introduced 
into  the  gas  it  is  extinguished;  nor  will  sulphur  burn  in  the  oxide. 


NITRIC  ACID,  NITROUS  ACID,  OXIDES  OF  NITROGEN  341 

If,  however,  a  bit  of  phosphorus  which  is  actively  burning  is  used, 
the  combustion  continues.  In  this  case  the  temperature  of  the 
flame,  which  is  higher,  is  sufficient  to  bring  about  the  decomposi- 
tion of  the  nitric  oxide. 

Nitric  oxide  takes  on  oxygen  with  great  readiness,  as  we  have 
seen.  When  the  two  gases  are  brought  into  contact  at  ordinary 
temperatures  nitrogen  dioxide,  a  brown  gas,  is  formed : 

2NO  +  O2  <=±  2NO2 

Nitric  oxide  unites  directly  with  other  elements  than  oxygen 
and  with  salts.  It  forms  nitrosyl  chloride  (363)  with  chlorine, 
2NO  +  C12  =  2NOC1,  and  nitrosyl  ferrous  sulphate,  FeSO4,NO, 
with  ferrous  sulphate.  All  these  compounds  decompose  more  or 
less  readily  into  their  constituents. 

381.  Nitrogen  Dioxide  and  Nitrogen  Tetroxide. — The  color  of 
nitrogen  dioxide  changes  markedly  with  change  in  temperature. 
At  150°  it  is  deep  brown,  and  as  the  temperature  falls  the  color 
decreases  in  intensity.  The  gas  condenses  to  a  bright  yellow  liquid 
at  26°;  the  color  of  this  slowly  fades  with  falling  temperature  and 
at  —9°  a  white  crystalline  solid  is  formed.  The  study  of  this 
phenomenon  has  led  to  the  conclusion  that  what  we  have  called 
nitrogen  dioxide  is  at  ordinary  temperatures  a  mixture  of  two 
compounds,  one  brown  and  one  colorless.  With  rise  in  temperature 
the  colorless  compound  is  slowly  converted  into  the  brown  oxide 
and  with  falling  temperature  the  reverse  phenomenon  takes  place; 
there  is  an  equilibrium,  markedly  affected  by  temperature,  which 
is  represented  by  the  following  equation: 

N2O4  <=*  2NO2  -  13,600  cal. 

The  brown  gas  is  nitrogen  dioxide  and  the  colorless  one  is  the 
polymer  nitrogen  tetroxide.  This  explanation  of  the  change  in 
color  has  been  arrived  at  by  determining  the  average  molecular 
weight  of  the  gas  at  different  temperatures. 

Nitrogen  dioxide  can  be  made  by  the  action  of  oxygen  or  air  on 
nitric  oxide  or  by  heating  nitrates  of  the  heavy  metals.  Copper 
nitrate,  as  we  have  seen,  yields  the  gas  when  heated  (367) : 

2Cu(NO3)2  =  2CuO  +  4N02  +  02 


342  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  gas  is  collected  by  passing  it  through  a  vessel  kept  cold  by  a 
mixture  of  ice  and  salt;  it  condenses  to  a  yellow  liquid.  Nitrogen 
dioxide  is  an  active  oxidizing  agent.  Fuming  nitric  acid  which 
contains  dissolved  nitrogen  dioxide  is  a  more  powerful  oxidizing 
agent  than  nitric  acid  itself. 

382.  Nitrogen   dioxide   can   be   detonated,  and  since  1  mole- 
cule furnishes  2  oxygen  atoms  which  can  be  utilized  for  oxidizing 
other  elements,  a  mixture  of  the  liquid  dioxide  and  compounds 
containing  carbon  and  hydrogen  is  a  powerful  explosive.     Per- 
fectly dry  nitrogen  dioxide  does  not  affect  benzene  or  naphthalene, 
but  when  the  mixture  is  detonated  an  exceedingly  violent  explosion 
results.     The  force  of  the  explosion  is  traceable  to  the  heat  lib- 
erated as  the  result  of  the  separation  of  the  oxide  into  its  elements 
and  that  produced  as  a  consequence  of  the  union  of  the  oxygen  with 
the  carbon  and  hydrogen  present  in  the  benzene  or  other  sub- 
stance used.     Only  gases  are  formed,  and  as  these  are  produced 
so  suddenly  that  the  heat  generated  cannot  escape,  a  tremendous 
pressure  is  produced.     Explosives  based  on  this  principle  were 
used  in  the  recent  war  in  depth  bombs. 

Nitrogen  tetroxide  reacts  with  cold  water  as  follows : 

N2O4  +  H2O  =  HNO2  +  HNO3 

If  the  gas  is  passed  into  hot  water  the  nitrous  acid  formed  decom- 
poses as  indicated  by  the  following  equation: 

3HNO2  =  HNO3  +  2NO  +  H2O 

By  combining  this  equation  with  the  one  given  above  we  have  the 
following : 

.    3N2O4  +  2H2O  =  4HNO3  +  2NO 

383.  Nitrous  Anhydride. — When  a  nitrite  is  treated  in  the  cold 
with  an  acid,  the  nitrous  acid  first  formed  spontaneously  decom- 
poses and  forms  a  brown  gas,  which  can  be  condensed  at   —21° 
to  a  blue  liquid  that  has  the  composition  represented  by  the  for- 
mula N2O3: 

2HN02  =  H2O  +  N2O3 

The  gas  reacts  with  solutions  of  bases  in  the  cold  and  forms  nitrites. 
For  these  reasons  nitrogen  trioxide  is  usually  called  nitrous  anhy- 


NITRIC  ACID,  NITROUS  ACID,  OXIDES  OF  NITROGEN  343 

dride.  The  compound  is  very  unstable  and  at  its  boiling-point 
begins  to  decompose  into  nitric  oxide  and  nitrogen  dioxide: 

N203  <=±  NO  +  N02 

At  ordinary  temperatures  the  dissociation  is  almost  complete, 
and,  consequently,  when  we  pass  the  gas  into  a  solution  of  a  base 
we  really  have  the  simultaneous  action  of  nitric  oxide  and  nitrogen 
dioxide  on  it. 

384.  Nitric  Anhydride. — Nitric  acid  does  not  break  down  spon- 
taneously into  its  anhydride  and  water,  but  if  it  is  treated  with  a 
powerful  dehydrating  agent  the  reaction  takes  place.  Phosphorus 
pentoxide,  which  can  be  used  for  this  purpose,  unites  with  the 
water  withdrawn  and  is  converted  into  metaphosphoric  acid: 

2HN03  +  P205  =  N205  +  2HP03 

When  the  mixture  is  cautiously  heated,  a  brown  gas  is  formed, 
which  condenses  to  a  liquid  that  boils  at  45°  and  freezes  to  a  white 
solid  at  30°.  The  anhydride  unites  with  water  and  forms  nitric 
acid,  and  decomposes  slowly  into  nitrogen  dioxide  and  oxygen: 

2N2O5  =  4N02  +  O2 

EXERCISES 

1 .  What  weight  of  Chile  saltpeter  containing  90  per  cent  NaNO2  is  required 
to  prepare  1  ton  of  concentrated  nitric  acid,  sp.  gr.   1.41,   provided   the 
process  is  carried  out  in  such  a  way  that  95  per  cent  of  the  nitric  acid  theo- 
retically obtainable  from  the  nitrate  is  obtained? 

2.  Write  equations  for  the  oxidation  by  nitric  acid  of  the  following:    (a) 
Na2SO3,    (6)  As2O3  to  As2O6,    (c)  P  to  P2O6;    (d)  CuHaOu  to  CO2  and  H2O, 
(e)  Fe  to  Fe(NO3)3. 

3.  (a)  Calculate  from  the  equation   representing  the  reaction  between 
hydrochloric  acid  and  nitric  acid  the  relation  between  the  weights  of  the 
acids  required  in  making  aqua  regia.     (6)  If  concentrated  hydrochloric  acid 
(sp.  gr.  1.2,  40  per  cent  HC1)  and  concentrated  nitric  acid  (sp.  gr.  1.4,  68  per 
cent  HNO3)  are  used,  what  is  the  relation  between  the  weights  of  these  acids 
required?     (c)  What  is  the  relation  between  the  volumes  of  the  concentrated 
acids  required? 

4.  Can  HNO3  be  used  to  prepare   (a)  HC1  from  NaCl,  and  (6)  CO,  from 
Na2C03? 

5.  When  HNO3  is  made  from  NO  prepared  from  N2  and  O2  a  large  amount 
of  energy  must  be  furnished  in  preparing  the  oxide.     When  the  acid  is  pre- 
pared by  oxidizing  NH3,  the  NO  first  formed  is  the  result  of  an  exothermic 


344  INORGANIC  CHEMISTRY  FOR  COLLEGES 

reaction;   the  union  of  N2  and  Ha  also  produces  heat.     At  what  stage  in  the 
second  synthesis  is  the  energy  furnished? 

6.  Why  is  concentrated  H2SO4  mixed  with  HNO3  in  making  cellulose 
nitrate? 

7.  Could  carbon  be  used  in  preparing  KNO2  from  KNO3?     Give  a  reason 
for  your  answer. 

8.  Explain  why  HNOa  is  a  more  active  oxidizing  agent  than  H2SO4. 

9.  How  would  you  distinguish   from  one  another  the    following:      (a) 
NH4NO3,    (6)  KNO2,  and   (c)  KNO3? 

10.  How  could  you  determine  whether  a  sample  of  N2O  contained  a  trace 
of  NO? 

11.  How  could  a  solution  of  HNO3  free  from  HNO>  be  made  from  NO? 
Should  the  water  used  be  hot  or  cold?     Why? 

12.  In  the  reversible  reaction  N2O4  <=^  2NO2  in  which  direction  is  the  reaction 
exothermic?      How  would  the  equilibrium  be  affected  by  rise  in  temperature? 

13.  How   would   increased   pressure   affect   the   dissociation   of   nitrogen 
tetroxide  when  it  is  heated? 

14.  Would  you  expect  that  N2O3  obeys    (a)  Boyle's  law  and    (6)  Charles' 
law?     Give  a  reason  for  your  answer  in  each  case. 

15.  Chile  saltpeter  from  which  nitric  acid  is  made  contains  a  small  percent- 
age of  sodium  chloride.     Name  two  impurities  that  are  introduced  into  the 
acid  as  a  result  of  the  presence  of  the  salt.     How  could  these  be  removed 
from  the  nitric  acid? 

16.  When  nitric  acid  is  prepared  it  has  a  brown  color  due  to  the  presence 
of  NOa,     How  could  this  color  be  removed? 


CHAPTER  XXIV 
THE  DETERMINATION  OF  ATOMIC  AND  MOLECULAR  WEIGHTS 

385.  Up  to  this  point  the  student  has  gained  a  sufficient 
familiarity  with  chemical  facts  to  make  a  consideration  of  the 
methods   of   determining   atomic   weights   profitable.     We   have 
already  seen  that  the  law  of  definite  proportions  and  the  law  of 
multiple  proportions  led  Dalton  to  study  the  problem  of  deter- 
mining the  relative  weights  of  atoms,  but  as  he  was  working  with 
compounds,  and,  therefore,  molecules,  and  had  no  way  of  deter- 
mining the  number  of  atoms  in  the  molecules  it  was  impossible 
for  him  to  arrive  at  figures  that  expressed  the  relative  weights  of 
the  atoms  themselves.     Dalton  found,  for  example,  that  water 
is  made  up  by  weight  of  1  part  of  hydrogen  and  8  parts  of  oyxgen. 
If  a  molecule  of  water  contains  1  atom  of  each  element  and  its 
formula  is  HO,  then  the  atomic  weight  of  oxygen  is  8  if  that  of 
hydrogen  is  1.     If,  however,  the  molecule  contains  2  atoms  of 
hydrogen  and  1  of  oxygen,  H2O,  the  atomic  weight  of  oxygen  is  16; 
and  other  atomic  proportions  lead  to  different  atomic  weights  for 
oxygen. 

386.  The  Law  of  Gay-Lussac.— In  1808  Gay-Lussac  published 
his  law  of  combining  volumes,  which  states  that  when  substances 
interact  in  the  gaseous  condition  the  volumes  of  the  reacting 
substances  and  those  of  the  products  formed  are  in  the  relation  of 
small  whole  numbers.     For  example,  2  volumes  of  hydrogen  unite 
with  1  volume  of  oxygen  and  form  2  volumes  of  steam;  1  of  hydro- 
gen and  1  of  chlorine  form  2  of  hydrogen  chloride,  etc. 

The  striking  siniDlicity  in  the  weight  relations  in  chemical 
reactions  summarized  in  the  law  of  multiple  proportions,  has  its 
counterpart  in  the  equally  striking  simplicity  in  the  volume  rela- 
tions observed  when  reaction  takes  place  between  substances  in 
the  gaseous  state.  We  have  seen  how  an  attempt  to  give  a  physical 

345 


346  INORGANIC  CHEMISTRY  FOR  COLLEGES 

interpretation  of  the  law  in  regard  to  weight  relations  led  to  an 
atomic  conception  of  matter  which  was  susceptible  of  experi- 
mental investigation  and,  therefore,  unlike  the  old  Greek  views; 
and  we  shall  now  see  how  the  law  summarizing  volume  relations, 
as  the  result  of  a  similar  endeavor,  furnished  the  key  to  the  prob- 
lem of  determining  the  relative  weights  of  molecules  and  atoms. 
The  law  of  multiple  proportions  has  to  do  with  the  relationship 
between  atoms;  the  law  of  Gay-Lussac  summarizes  the  relation- 
ship between  molecules.  And  we  have  just  seen  that  Dalton's 
attempt  to  determine  atomic  weights  was  futile  because  he  had  no 
way  of  determining  the  number  of  atoms  in  molecules. 

The  significance  of  Gay-Lussac's  law  was  first  seen  by  Avo- 
gadro,  an  Italian,  in  1811,  but  the  important  conclusions  which  he 
drew  from  it  were  not  appreciated  until  later,  when  Cannizzaro, 
in  1860,  demonstrated  how  Avogadro's  conception  furnished  a 
satisfactory  method  upon  which  to  base  the  determination  of 
atomic  weights.  For  years  chemists  had  been  changing  from  one 
system  to  another,  and  a  number  of  these  were  in  use;  as  a  conse- 
quence, one  chemist  wrote  the  formula  H^O  for  water  and  another 
HO.  The  system  of  atomic  weights  used  to-day  is  based  on  the 
method  of  determination  outlined  by  Cannizzaro,  who  was  able  to 
clear  up  the  difficulties  existing  because  he  possessed  the  power  to 
marshal  the  facts  in  such  a  simple  and  logical  way  that  their  sig- 
nificance was  understood. 

387.  The  Law  of  Avogadro. — Avogadro  clearly  differentiated 
atoms  from  molecules — the  molecule  was  the  smallest  physical 
unit  of  a  substance  and  was  made  up  of  atoms.  And  on  account 
of  the  simplicity  of  the  volume  relations  involved  when  molecules 
interact  he  emphasized  the  necessity  of  considering  these  relations 
because  they  were  without  doubt  related  to  the  volumes  of  the 
molecules  themselves. 

Let  us  see  how  a  study  of  the  volumes  occupied  by  gaseous 
substances  helps  in  the  problem  which  baffled  Dalton.  In  the 
selection  of  a  standard  of  atomic  weights  the  logical  thing  to  do  is 
to  select  the  lightest  element  known  and  assign  to  it  the  atomic 
weight  1;  this  was  done  and  hydrogen  was  selected  (66).  It 
would  be  interesting  at  the  outset  to  compare  the  volumes  occu- 
pied by  a  number  of  gaseous  substances,  and  to  have  a  basis  for 
the  comparison  we  shall  take  the  amount  of  each  gas  which  con- 


DETERMINATION  OF  ATOMIC  AND  MOLECULAR  WEIGHTS  347 


tains  1  gram  of  hydrogen.     These  volumes  are  recorded  in  the 
following  table: 


Element  combined  with 
hydrogen     

Cl 

o 

N 

c 

S 

P 

Si 

H 

Volume  in  liters  at  0°  and 
760   mm.    containing    1 
gram  of  hydrogen  
Relative    volumes    with 
largest  as  unity  

22.4 
1. 

11.2 
i 

7.5 
A 

5.5 
| 

11.2 

i 

7.5 
\ 

5.5 
i 

11.2 

A 

In  the  last  column  is  given  hydrogen  gas.  The  simple  relation 
that  exists  between  the  volumes  of  the  gases  listed  which  contain  1 
gram  of  hydrogen  is  striking;  and  we  would  obtain  similar  results 
if  other  hydrogen  compounds  were  studied.  A  natural  question 
is — What  is  the  physical  basis  for  this?  It  is  hardly  possible  that 
the  molecules  of  the  different  gases  all  contain  the  same  number  of 
hydrogen  atoms.  What  would  be  the  result  if  they  differed  in  this 
respect?  If  we  prepare  from  the  same  amount  of  hydrogen  in 
each  case  molecules  which  contain  respectively  1,  2,  3,  and  4  atoms 
of  hydrogen,  the  number  of  molecules  formed  in  the  second  case  will 
be  one-half  as  many  as  those  formed  in  the  first  case;  the  number 
formed  in  the  third  case  will  be  one-third  as  many  as  in  the  first 
case;  and  the  number  formed  in  the  fourth  case  one-fourth  as  many. 
This  fact  when  considered  along  with  the  fact  that  the  volumes  of 
the  gases  formed  from  the  fixed  weight  of  hydrogen  are  in  the 
relation  of  1  to  i,  J,  and  J  leads  to  the  reasonable  assumption  that 
the  volume  of  a  gas  is  determined  by  the  number  of  molecules  it 
contains,  whatever  the  molecule  may  be.  This  conclusion  is 
known  as  Avogadro's  Law  and  is  usually  stated  as  follows :  Equal 
volumes  of  all  gases  at  the  same  temperature  and  pressure  contain 
the  same  number  of  molecules. 

For  years  this  remarkable  conclusion  of  Avogadro  was  known 
as  a  hypothesis  for  no  one  was  able  to  count  the  number  of  mole- 
cules of  a  gas;  and  it  seemed  impossible  that  physical  experimen- 
tation could  ever  reach  such  a  point  of  refinement  that  it  could 
deal  with  particles  so  small  that  billions  of  billions  of  them  could 
be  contained  in  a  space  the  size  of  a  drop  of  water.  The  hypothesis 
became,  however,  the  basis  of  the  determination  of  atomic  weights 


348  INORGANIC  CHEMISTRY  FOR  COLLEGES 

and,  therefore,  had  an  important  significance;  it  was,  accordingly, 
known  later  as  Avogadro's  rule.  Recent  work  in  physics  has  made 
it  possible  to  count  molecules,  and  as  a  number  of  independent 
methods  all  lead  to  the  same  conclusion  and  confirm  the  hypoth- 
esis, it  is  now  known  as  a  law.  One  gram  of  hydrogen  contains 
3  X  1023  molecules.1 

388.  The  Determination  of  Atomic  Weights. — Let  us  now  see 
how  the  fact  that  equal  volumes  of  gases  contain  the  same  number 
of  molecules  can  be  used  in  building  up  a  system  of  atomic  weights. 
We  will  take  as  our  standard  hydrogen — the  lightest  known  sub- 
stance— and  call  its  atomic  weight  1.  The  volumes  of  as  many 
compounds  of  hydrogen  as  possible  are  determined,  selecting  in 
each  case  the  volume  which  contains  1  gram  of  hydrogen.  The 
results  in  a  number  of  cases  are  given  in  the  table  already  dis- 
cussed (page  347).  It  is  evident  that  we  shall  get  the  largest 
volume  in  the  case  of  the  gas  that  has  but  1  atom  in  the  molecule 
for  in  this  gas  we  shall  have  the  largest  number  of  molecules. 
Accordingly,  in  the  compound  which  occupies  the  largest  volume 
we  assume  that  1  atom  of  hydrogen  is  present  in  the  molecule. 
This  assumption  is  justified  when  we  select  the  compound  contain- 
ing chlorine,  hydrochloric  acid,  for  there  is  no  gas  known  which  has 
a  larger  volume  than  22.4  liters  when  the  amount  of  gas  selected 
is  that  which  contains  1  gram  of  hydrogen.  These  considerations 
lead  us  to  the  view  that  the  molecule  of  hydrochloric  acid  contains 
1  hydrogen  atom.  We  can  determine  the  number  of  chlorine 
atoms  in  the  molecule  in  the  same  way  by  examining  the  volumes 
of  chlorine  compounds  all  of  which  contain  the  same  weight  of  the 
element.  We  find  in  this  case  that  none  of  these  yields  a  greater 
volume  than  the  compound  containing  hydrogen,  and  we  con- 
clude, therefore,  that  hydrochloric  acid  contains  1  chlorine  atom. 
If  there  is  1  atom  of  hydrogen  and  1  atom  of  chlorine  in  hydro- 
chloric acid  its  formula  is  HC1.  We  find  by  experiment  that  1 
gram  of  hydrogen  unites  with  35.5  grams  of  chlorine  and  forms  36.5 
grams  of  hydrochloric  acid;  the  atomic  weight  of  chlorine  is, 
therefore,  35.5,  and  the  molecular  weight  of  hydrochloric  acid  is 
36.5.  We  also  find  by  experiment  that  36.5  grams  of  hydrochloric 
acid — 1  gram-molecular-weight  or  mol — occupies  22.4  liters. 

1  This  is  a  short  way  of  expressing  the  number  made  up  of  3  followed  by 
23  zeros;  it  is  300,000  billion  billions. 


DETERMINATION  OF  ATOMIC  AND  MOLECULAR  WEIGHTS     349 

We  have  thus  a  definite  experimental  basis  founded  on  fact  for 
the  determination  of  the  atomic  weight  of  chlorine  and  the  formula 
of  hydrochloric  acid;  and  the  same  method  can  be  applied  in  other 
cases.  The  only  thing  that  would  lead  to  a  change  in  the  atomic 
weight  of  chlorine  would  be  the  discovery  of  a  compound  which 
occupied  as  a  gas  more  than  22.4  liters,  when  the  weight  of  it  con- 
taining 1  gram  of  hydrogen  was  used;  for  the  gas  which  occupies 
the  largest  volume  has  1  atom,  and  in  this  case  there  would  be  more 
than  1  hydrogen  atom  in  hydrochloric  acid. 

The  method  used  in  determining  atomic  weights  can  be  further 
emphasized  by  applying  it  to  oxygen.  The  same  reasoning  is  used. 
The  volumes  of  gaseous  compounds  containing  oxygen  are  deter- 
mined, using  in  each  case  the  amounts  of  the  several  substances 
which  contain  the  same  weight  of  oxygen.  Results  which  have 
been  obtained  in  a  number  of  compounds  are  given  in  the  following 
table: 


Element  combined  with 
oxygen  

H  as 

C  as 

Gas 

Sas 

Sas 

Nas 

Nas 

Vol.  of  substance  at  0° 
and  760  mm.  contain- 
ing 1  gram  oxygen  

H2O 
1.4 

CO 
1.4 

CO2 
0.7 

SO2 
0.7 

SO3 
0.48 

NO 
1.4 

NO2 
0.7 

Relative  volumes  with 
largest  as  unity 

1 

1 

| 

| 

i 

1 

i 

From  these  results  we  draw  the  conclusion  that  water,  carbon 
monoxide,  and  nitric  oxide  each  contains  1  atom  of  oxygen.  An 
examination  of  the  table  on  page  347  will  give  us  information  as  to 
the  number  of  hydrogen  atoms  in  water.  The  volume  of  the 
water-vapor  which  contains  1  gram  of  hydrogen  is  one-half  the 
volume  of  the  hydrochloric  acid  containing  this  weight  of  hydrogen, 
and,  therefore,  there  are  twice  as  many  atoms  of  hydrogen  in  water 
as  in  hydrochloric  acid,  and  its  formula  is  H^O.  Since  18  grams 
of  water  contain  2  grams  of  hydrogen  and  16  grams  of  oxygen,  the 
atomic  weight  of  oxygen  is  16  when  hydrogen  is  1. 

389.  The  Determination  of  Formulas. — Having  arrived  in  this 
way  at  the  atomic  weights  of  the  elements  we  can  readily  deter- 
mine the  formula  of  any  compound.  Avogadro's  law  tells  us  that 


350  INORGANIC  CHEMISTRY  FOR  COLLEGES 

equal  volumes  of  all  gases  contain  the  same  number  of  molecules; 
as  a  consequence,  the  weights  of  equal  volumes  of  any  two  gases 
are  in  the  same  relation  as  the  weights  of  the  individual  molecules 
of  these  gases.  We  have  seen  that  22.4  liters  of  hydrochloric 
acid  weigh  36.5  grams  and  that  the  molecular  weight  of  hydro- 
chloric acid  is  36.5.  If,  therefore,  we  determine  the  weight  in 
grams  of  22.4  liters  of  any  gas  the  number  obtained  will  be  the 
molecular  weight  of  the  gas.  The  volume  22.4  liters  is,  there- 
fore, one  of  the  greatest  significance  in  chemistry;  for  it  is  the 
volume  of  a  gram-molecular-weight  of  all  gases.  We  can  deter- 
mine, therefore,  the  molecular  weight  of  any  gas  by  weighing  a 
sample  of  known  volume  and  calculating  from  the  result  the  weight 
of  22.4  liters — the  number  obtained  is  the  molecular  weight  of  the 
gas. 

In  order  to  write  the  formula  of  a  substance  we  must  know,  in 
addition  to  the  weight  of  its  molecule,  the  proportion  by  weight  of 
the  elements  of  which  it  is  composed,  and  the  atomic  weights  of 
these  elements.  An  example  will  show  how  this  can  be  done. 
A  substance  was  found  upon  analysis  to  contain  30.43  per  cent 
nitrogen  and  69.57  per  cent  oxygen;  the  weight  of  22.4  liters  at 
0°  and  760  mm.  was  46  grams.  What  is  its  formula?  The 
molecular  weight  is  46;  30.43  per  cent  of  this  is  nitrogen,  there- 
fore, 46  X  .3043  =  14  is  the  weight  of  the  nitrogen  in  the  mole- 
cule. The  weight  of  the  oxygen  is  46  X  .6957  =  32.  Since 
an  atom  of  nitrogen  weighs  14  and  1  of  oxygen  weighs  16  the  for- 
mula is,  evidently,  NO2.  The  formula  of  any  compound  can  be 
determined  in  this  way.  The  weight  of  22.4  liters  at  0°  and  760 
mm.  of  the  substance  in  the  form  of  a  gas  is  multiplied  in  turn  by 
the  percentage  of  each  element  present,  and  the  result  in  each  case 
divided  by  the  atomic  weight  of  the  element.  The  numbers 
obtained  are  the  numbers  of  the  several  atoms  present  in  the 
molecule. 

390.  The  formulas  of  elementary  gases  can  be  determined  in 
the  same  way.  For  example,  22.4  liters  of  hydrogen  weigh  2 
grams.  Since  the  atomic  weight  of  hydrogen  is  1  the  formula  of 
hydrogen  gas  is  Eb. 

It  is,  of  course,  not  necessary  for  the  substance  to  exist  as  a 
gas  at  0°  in  order  to  get  its  molecular  weight  in  the  gaseous  con- 
dition. The  volume  of  a  given  weight  is  determined  at  any  con- 


DETERMINATION  OF  ATOMIC  AND  MOLECULAR  WEIGHTS     351 

venient  temperature  and  pressure  and  a  calculation  is  made  by 
applying  the  gas  laws  to  determine  what  the  volume  would  be  at  0° 
and  760  mm.  In  this  way  the  molecular  weight  of  phosphorus  in 
the  gaseous  condition  has  been  determined  although  the  element 
is  a  solid  at  ordinary  temperatures.  Since  the  molecular  weight 
was  found  to  be  124  and  the  atomic  weight  is  31  the  molecule  in 
the  gaseous  condition  at  the  temperature  used  is  P4. 

391.  Atomic  Weights  Based  on  Oxygen  as  Standard. — The 
system  of  atomic  weights  which  has  been  explained  is  based  on 
the  hydrogen  atom  to  which  was  assigned  the  value  1,  and  round 
numbers  have  been  used  in  order  to  simplify  the  discussion.  The 
most  careful  analyses  of  water  that  have  been  made  lead  to  the 
conclusion  that  2  grams  of  hydrogen  unite  with  15.88  grams  of 
oxygen  and,  as  a  consequence,  the  atomic  weight  of  oxygen  is 
15.88  if  that  of  hydrogen  is  1.  Using  this  value  the  atomic  weight 
of  other  elements  can  be  determined;  for  example,  the  weight  of 
copper  that  unites  with  15.88  grams  of  oxygen  to  form  copper 
oxide,  CuO,  is  the  atomic  weight  of  copper,  namely  63.09. 

The  atomic  weights  of  a  large  number  of  elements  were  found 
by  determining  the  relative  weights  of  the  element  and  oxygen  in 
their  oxides;  and  the  values  of  the  atomic  weights  were  calculated 
with  the  aid  of  the  atomic  weight  of  oxygen.  The  value  of  the 
latter,  as  we  have  seen,  is  deduced  from  the  results  of  the  analysis 
of  water,  and  the  actual  numbers  used  for  the  atomic  weights  of 
all  elements  calculated  from  the  analysis  of  their  oxides  is  deter- 
mined, therefore,  by  the  figures  obtained  in  the  study  of  the  com- 
position of  water.  As  methods  of  analysis  increased  in  accuracy — 
as  balances  were  improved,  for  example — the  results  of  the  analysis 
of  water  gave  different  and  more  accurate  figures,  and  the  atomic 
weight  assigned  to  oxygen  kept  changing.  For  a  long  time  15.79 
was  accepted  as  the  atomic  weight  of  oxygen  and  the  atomic 
weight  of  other  elements  calculated  from  this  figure.  Then  more 
accurate  work  showed  that  2  grams  of  hydrogen  united  not  with 
15.79  grams  of  oxygen,  but  with  15.88  grams;  and  15.88  was  taken 
as  the  atomic  weight  of  the  element.  As  a  result,  the  atomic 
weights  of  all  elements  which  were  determined  by  analyzing  oxides 
had  to  be  recalculated  and  new  figures  were  obtained.  Every 
time  the  atomic  weight  of  oxygen  was  changed,  as  a  result  of  a 
more  accurate  determination  of  the  ratio  between  the  weights  of 


352  INORGANIC  CHEMISTRY  FOR  COLLEGES 

hydrogen  and  oxygen  that  unite,  all  atomic  weights  had  to  be  recal- 
culated. This  was  an  unfortunate  state  of  affairs  and  resulted 
largely  from  the  fact  that  the  accurate  determination  of  the  ratio 
between  the  weights  of  hydrogen  and  oxygen  that  unite  to  form 
water  is  one  of  the  most  difficult  of  analytical  processes,  because 
hydrogen  is  the  lightest  substance  known  and  the  weighing  and 
measuring  of  large  volumes  of  a  gas  are  exceedingly  difficult 
operations.  To  see  what  could  be  done  to  avoid  this  constant 
changing  of  atomic  weights  an  international  commission  of  chem- 
ists was  appointed  to  devise  a  plan  to  be  followed.  This  resulted  in 
taking  as  the  standard  oxygen  and  calling  its  atomic  weight  16. 
As  a  result,  the  atomic  weight  of  hydrogen  became  1.008,  and  any 
changes  in  the  oxygen-hydrogen  ratio  affected  the  atomic  weight 
of  hydrogen  alone.  The  value  accepted  for  the  atomic  weight 
of  copper,  for  example,  will  be  changed  only  when  more  accurate 
analyses  of  copper  compounds  lead  to  different  results  from  those 
obtained  in  the  past. 

392.  The  Determination  of  Atomic  Weights  by  Analysis. — 
It  has  been  pointed  out  that  the  accurate  measurement  of  gases  is 
difficult,  and,  as  a  consequence,  the  values  of  the  atomic  weights 
obtained  in  this  way  are  not  the  most  accurate  that  can  be  arrived 
at.     We  determine,  for  example,  in  the  way  described,  from  gases 
containing  carbon,  that  its  atomic  weight  is  approximately  12, 
but  we  arrive  at  a  more  accurate  figure  by  using  solid  or  liquid 
compounds  containing  the  element.     A  method  that  could  be  used 
will  illustrate  the  procedure.     A  carefully  weighed  amount  of 
carbon  could  be  burned  in  a  stream  of  oxygen  and  the  gas  formed, 
CO2,  absorbed  by  sodium  hydroxide.     The  increase  in  weight  of 
the  vessel  containing  the  sodium  hydroxide  gives  the  weight  of 
carbon  dioxide  produced.     The  difference  between  this  weight  and 
the  weight  of  carbon  burned  is  the  weight  of  the  oxygen.     From 
the  weight  of  carbon  and  the  weight  of  oxygen  and  knowing  the 
formula  of  carbon  dioxide  and  the  atomic  weight  of  oxygen,  the 
atomic  weight  of  carbon  can  be  calculated. 

393.  The  Law  of  Dulong  and  Petit. — The  system  of  determining 
atomic  weights  which  has  been  described  is  based  on  conclusions 
drawn  from  a  study  of  the  volumes  of  gaseous  compounds  of  the 
elements.     If  such  compounds  do  not  exist,  we  must  employ  some 
other  method  which  yields  results  that  fit  into  the  system  in  use. 


DETERMINATION  OF  ATOMIC  AND  MOLECULAR  WEIGHTS     353 

In  general,  the  acid-forming  elements  yield  compounds  that  are 
gases  or  can  be  readily  vaporized,  while  most  metals  do  not  form 
such  compounds.  If  a  method  could  be  found  to  determine  the 
atomic  weights  of  metals  which  did  not  involve  the  use  of  gases  it 
could  be  linked  up  with  the  method  that  has  been  described, 
because  some  metals,  such  as  mercury,  do  form  compounds  which 
can  be  vaporized.  A  determination  of  the  atomic  weight  of  such  a 
metal  by  the  two  methods  should  give  the  same  result  if  the  one 
applied  to  metallic  atoms  is  to  be  used;  for  it  is  necessary  to  have 
the  same  standard  for  all  atomic  weights. 

All  gases  have  the  same  molecular  volume — that  is,  a  molecule 
of  any  gas  occupies  the  same  space  as  a  molecule  of  any  other  gas; 
and  we  have  used  this  fact  to  build  up  a  system  of  atomic  weights. 
There  is  no  such  generalization  in  the  case  of  liquids  and  solids, 
and  some  other  property  than  the  relation  between  weight  and 
volume  must  be  used  in  deciding  upon  the  atomic  weights  of  metals. 
Dulong  and  Petit,  two  French  chemists,  investigated  the  specific 
heats  of  solid  elements  and  discovered  an  important  generalization 
which  was  announced  in  1819.  It  will  be  recalled  that  the  specific 
heat  of  a  substance  is  the  number  of  calories  required  to  raise  the 
temperature  of  1  gram  of  it  1  degree.  The  specific  heats  of  elements 
vary  widely;  a  few  values  for  metals  are  as  follows:  iron,  0.112; 
zinc,  0.093;  gold,  0.032.  It  was  noticed  that  in  a  series  of  metals 
as  the  atomic  weight  increased  the  specific  heat  decreased;  and  the 
investigators  were  led  to  multiply  these  values  in  the  case  of  each 
element.  The  products  obtained  were  approximately  the  same 
number.  For  example,  iron,  atomic  weight  56,  specific  heat  0.112, 
56  X  0.112  =  6.3;  gold  197  X  0.032  =  6.3;  zinc  65.4  X  0.093  =  6.1. 
The  average  value  of  the  product  in  the  case  of  metals  is  6.4.  Such 
facts  as  these  led  to  what  is  called  Dulong  and  Petit's  law,  which 
states  that  the  product  obtained  by  multiplying  the  specific  heat 
of  a  metal  by  its  atomic  weight  is  a  constant  in  all  cases.  The 
physical  significance  of  this  is  clear.  The  product  obtained  in 
each  case,  approximately  6.4  calories,  is  the  heat  required  to  raise 
1  gram-atomic-weight  of  the  metal  1  degree.  So  the  law  can  be 
stated  in  the  following  form:  "  The  atoms  of  all  metals  have  the 
same  heat  capacity."  This  generalization  will,  accordingly,  give 
us  a  means  of  determining  the  relative  atomic  weights ;  all  that  is 
needed  is  a  standard,  and  this  must  be  the  same  as  that  used  in 


354  INORGANIC  CHEMISTRY  FOR  COLLEGES 

the  case  of  the  gaseous  elements.  It  is  well  to  point  out  here  again 
that  the  system  of  determining  the  atomic  weights  of  elements 
that  exist  in  gaseous  compounds  is  based  on  the  fact  that  all 
molecules  occupy  the  same  volume;  and  we  shall  see  that  the 
system  for  obtaining  the  atomic  weights  of  solid  elements  is  based 
upon  the  fact  that  the  atoms  of  all  metals  have  the  same  capacity 
for  heat. 

In  order  to  determine  the  atomic  weight  of  a  metal  all  that  is 
necessary  is  to  determine  its  specific  heat  and  make  use  of  the  fact 
that 

atomic  weight  X  specific  heat  =  6.4 

If  the  specific  heat  of  a  newly  discovered  metal  were  determined  to 
be  0.04  its  atomic  weight  would  be  found  as  follows: 

at.  wt.  X  0.04  =  6.4 
at.  wt.  =  160 

The  application  of  the  law  of  Dulong  and  Petit  yields  approxi- 
mate values  only  for  the  atomic  weights  because  the  product  of 
the  specific  heat  and  the  atomic  weight  is  not  exactly  the  same 
number  in  all  cases.  The  reasons  for  these  variations  were  un- 
known for  a  long  time.  Recently,  however,  the  advance  in  our 
knowledge  of  energy  and  its  transfer  has  helped  in  interpreting  the 
deviations  from  the  law  of  Dulong  and  Petit.  Extension  of  our 
knowledge  of  matter  has  frequently  resulted  from  a  search  for  a 
reason  for  the  fact  that  the  so-called  constant  in  a  law  varies  slightly 
when  we  apply  the  law  to  specific  cases.  The  deviations  are  often 
very  small,  and  they  are  discovered  only  when  measurements  are 
made  with  the  greatest  care — a  fact  which  emphasizes  again  the 
value  of  the  most  accurate  work  possible  in  studying  the  quantity 
relationships  in  physical  and  chemical  phenomena.  By  the  use 
of  the  law  of  Dulong  and  Petit  the  approximate  value  of  the  atomic 
weight  of  a  metal  can  be  determined ;  the  accurate  value  is  obtained 
by  the  analysis  of  a  compound  containing  the  element.  For 
example,  we  calculate  from  the  specific  heat  of  copper  that  its 
atomic  weight  is  approximately  63.  By  a  careful  analysis  of 
copper  oxide  we  find  that  63.57  grams  of  copper  are  combined  with 
16  grams  of  oxygen;  the  atomic  weight  of  the  element  is,  therefore, 
63.57. 


DETERMINATION  OF  ATOMIC  AND  MOLECULAR  WEIGHTS     355 

The  law  of  Dulong  and  Petit  gives  us  a  way  to  determine  the 
relative  atomic  weights  of  the  metals,  but  it  does  not  fix  a  standard. 
In  order  to  get  the  constant  to  be  used  in  calculating  the  atomic 
weights  of  elements  from  their  specific  heat,  we  must  multiply  the 
specific  heat  of  some  element  by  its  atomic  weight — we  must  have  a 
standard.  Fortunately  some  metals  form  volatile  compounds  and 
we  can  determine  the  atomic  weights  of  such  metals  with  the  aid 
of  ttn  method  applied  to  gases  and  obtain  values  based  on  the 
standard — oxygen  16 — used  in  the  case  of  the  non-metallic  ele- 
ments. 

The  law  of  Dulong  and  Petit,  we  have  seen,  leads  to  the  con- 
clusion that  the  same  amount  of  heat  is  required  to  cause  the  same 
rise  in  temperature  of  the  atoms  of  all  metals.  This  is  a  most 
remarkable  fact;  we  would  not  expect  that  the  same  amount  of 
heat  would  produce  approximately  the  same  change  in  temperature 
in  the  case  of  a  light  atom  like  sodium,  which  weighs  23,  and  a 
heavy  atom  like  lead,  which  weighs  207.  As  a  result  of  the  appli- 
cation of  recent  theories  in  physics  to  the  study  of  atoms,  new 
light  is  being  thrown  on  this  important  fact.  In  the  whole  field 
of  chemistry  there  is  opportunity  for  research;  we  are  just  at  the 
threshold  of  the  science. 

394.  The  Importance  of  Accurate  Atomic  Weights. — The 
accurate  determination  of  the  atomic  weights  of  the  elements  is 
one  of  the  most  important  problems  in  chemistry,  and  great 
ingenuity  has  been  shown  in  increasing  the  accuracy  of  methods  of 
analysis  used  in  their  determinations.  The  atom  is  the  chemical 
unit  of  matter,  and  the  more  certain  our  knowledge  is  of  its  prop- 
erties the  more  fully  we  understand  the  changes  in  matter  and  the 
better  we  can  make  use  of  them.  The  weight  of  an  atom  is  but 
one  of  its  properties,  but  it  is  an  important  one;  one  of  the  most 
fundamental  generalizations  of  chemistry — the  periodic  law — is 
based  on  the  atomic  weights,  and  this  could  have  been  discovered 
only  after  reasonably  correct  values  of  the  atomic  weights  were 
known.  It  is  highly  probable  that  as  the  knowledge  of  the  true 
atomic  weights  becomes  more  exact  other  generalizations  will  be 
apparent,  and  causes  for  phenomena  not  yet  understood  will 
become  evident. 

The  determinations  of  the  atomic  weights  of  some  of  the  ele- 
ments are  among  the  most  accurate  measurements  which  have  ever 


356  INORGANIC  CHEMISTRY  FOR  COLLEGES 

been  made  in  science.  A  number  of  chemists  specially  equipped 
for  this  kind  of  work,  which  requires  physical  dexterity,  infinite 
patience,  the  ability  to  see  hidden  sources  of  error,  and  the  ingenu- 
ity necessary  to  invent  new  and  more  accurate  methods  of  measure- 
ment, have  obtained  the  admirable  results  summarized  in  a  table 
of  atomic  weights.  Stas,  a  Belgian  chemist  (1813-1891),  set  an 
example  to  the  chemical  world  by  his  work  in  determining  atomic 
weights.  His  results  were  accepted  as  the  standard  of  accuracy 
until  revised  later  by  Professor  T.  W.  Richards  of  Harvard  Uni- 
versity, who,  with  his  co-workers,  has  obtained  values  for  the 
atomic  weights  of  the  more  important  elements,  which  surpass  in 
accuracy  all  previous  determinations.  Professor  Edward  Morley, 
formerly  of  Adelbert  College,  spent  a  number  of  years  studying 
the  hydrogen-oxygen  ratio,  and  his  results  are  universally  accepted 
as  the  most  accurate  that  have  been  obtained  for  this  constant. 

It  is  necessary  that  chemists  the  world  over  agree  to  use  the 
same  values  for  the  atomic  weights,  and  as  these  numbers  are  con- 
stantly being  revised,  an  international  committee  formed  of  rep- 
resentatives from  the  chemical  societies  of  the  world  studies  the 
new  work  in  this  field  and  publishes  each  year  a  table  of  atomic 
weights.  Before  there  was  an  international  agreement  as  to  the 
value  of  atomic  weights  difficulties  arose  which  in  some  cases 
affected  commercial  relations.  At  one  time,  for  example,  the 
accepted  atomic  weight  for  chromium  was  not  the  same  in  Eng- 
land and  America.  In  calculating  the  amount  of  chromium  in 
an  ore  from  its  analysis,  it  is  necessary  to  use  the  atomic  weight 
of  the  element,  and  the  amount  found,  and  consequently  the  value 
of  the  ore  will  be  different  if  different  atomic  weights  are  used. 
It  is  necessary,  therefore,  in  commercial  transactions  of  this  kind 
to  agree  on  atomic  weights,  and  the  values  selected  by  the  inter- 
national commission  are  accepted  and  used  in  all  calculations. 

EXERCISES 

1.  A  number  of  compounds  gave  the  following  results.     Calculate  the 
formula  in  each  case:    (a)  91.30  per  cent  P  and  8.65  per  cent  H2;    1  liter 
weighed  1.52  g.     (6)  87.59  per  cent  Si  and  12.45  per  cent  H2;  1  liter  weighed 
1.44  grams,     (c)  80.20  per  cent  C  and  20.1  per  cent  H2;   1  liter  weighed  1.33 
grams,     (d)  92.36  per  cent  C  and  7.8  per  cent  H2;    1  liter  weighed   1.16 
grams. 

2.  An  experiment  showed  that  when  1.2714- grams  of  copper  were  heated 


DETERMINA  TION  OF  A  TOM  1C  AND  MOLECULAR  WEIGHTS    3  57 

in  a  stream  of  air  1.5920  grams  of  CuO  were  formed.     Calculate  the  atomic 
weight  of  Cu,  assuming  the  formula  of  the  oxide  to  be  CuO. 

3.  In  a  determination  of  the  atomic  weight  of  bromine  2.6970  grams 
of  silver  were  dissolved  in  HNO3  and  precipitated  as  AgBr;    the  weight 
of  the  latter  obtained  was  4.6950  grams.     Calculate  the  atomic  weight  of 
bromine,  using  107.88  as  the  atomic  weight  of  silver. 

4.  In  determining  the  atomic  weight  of  sodium   1.2356  grams  of  pure 
NaCl  were  dissolved  in  water  and  treated  with  a  solution  of  silver  nitrate. 
The  silver  chloride  obtained  weighed  3.0274  grams.     Taking  107.88  as  the 
atomic  weight  of  silver  and  35.46  as  that  of  chlorine  calculate  the  atomic 
weight  of  sodium. 

5.  The  specific  heat  of  a  metal  was  found  to  be  0.033.     When  5.0150 
grams  of  the  metal  were  converted  into  an  oxide  the  latter  weighed  5.4150 
grams.     Calculate  the  atomic  weight  of  the  element. 

6.  The  specific  heat  of  a  metal  was  found  to  be  0.031.     When  4.1440 
grams  of  the  metal  were  converted  into  its  chloride  the  latter  weighed  5.5624 
grams.     Calculate  the  atomic  weight  of  the  metal,  using  35.46  as  the  atomic 
weight  of  chlorine. 


CHAPTER  XXV 
THE  PERIODIC  LAW 

395.  The  elements  of  which  the  world  is  made  up  show  a  great 
variety  of  physical  and  chemical  properties.  Some  are  light  gases 
and  some  are  dense  solids;  some,  like  sulphur,  are  exceedingly 
poor  conductors  of  heat  and  electricity,,  and  others,  like  silver,  are 
excellent  conductors.  There  is  just  as  wide  a  range  in  chemical 
properties ;  we  have  active  metals  that  form  strong  bases,  of  which 
sodium  is  an  example,  and  equally  active  non-metals,  like  chlorine, 
that  yield  strong  acids.  Other  elements  are  characterized  by 
being  excessively  inert;  nitrogen  and  carbon  show  little  or  no 
activity  at  ordinary  temperatures.  Elements  exist  that  resemble 
each  other  strikingly  in  chemical  properties.  Calcium,  for  exam- 
ple, forms  a  series  of  compounds  the  chemical  behavior  of  which  is 
markedly  like  that  of  the  compounds  of  magnesium  of  analogous 
composition.  The  question  before  the  chemist  has  always  been 
what  is  the  fundamental  cause  underlying  these  differences  and 
similarities. 

After  Dalton  had  proposed  the  atomic  theory  and  the  atomic 
weights  of  a  number  of  elements  had  been  determined,  it  was 
noticed  that  there  was  some  relation  between  these  numbers  and 
the  properties  of  the  elements.  The  elements  chlorine,  bromine, 
and  iodine  resemble  one  another  markedly  in  chemical  properties; 
they  form  with  other  elements  compounds  of  analogous  composi- 
tion, which  resemble  one  another  in  their  chemical  behavior. 
When  the  physical  properties  of  these  elements  and  their  com- 
pounds are  studied  it  is  found  that  there  is  a  progressive  change  as 
we  pass  from  chlorine  to  bromine  to  iodine.  For  example,  chlorine 
is  a  yellow  gas,  bromine  a  red  liquid,  and  iodine  a  black  solid. 
The  solubilities  of  the  compounds  of  the  formulas  NaCl,  NaBr, 
and  Nal,  are,  respectively,  35.86,  88.76,  and  177.9  grams  in  100  c.c. 

358 


THE  PERIODIC  LAW  359 

of  water  at  18°.  On  account  of  this  close  relationship  these  ele- 
ments were  said  to  belong  to  a  chemical  family.  The  atomic 
weights  of  chlorine,  bromine,  and  iodine  are,  respectively,  35.46, 
79.92,  and  126.92.  The  increase  in  these  values  is  progressive, 
and  it  is  a  remarkable  fact  that  the  atomic  weight  of  bromine  is  not 
far  from  the  mean  of  the  values  of  the  other  two,  which  is  81.19. 
These  facts  are  of  interest,  but  become  of  great  significance  when 
considered  along  with  the  fact  that  other  so-called  families  of 
elements  exist,  and  that  in  these  there  is  the  same  numerical 
relationship  between  the  atomic  weights  of  the  members  of  any 
family. 

It  was  clear  that  there  was  some  connection  between  the  prop- 
erties of  elements  and  their  atomic  weights,  and  a  number  of 
attempts,  more  or  less  successful,  were  made  to  correlate  these 
facts.  It  was  some  time  after  they  were  known,  however,  before 
the  real  relationship  was  discerned.  Mendelejeff,  a  Russian  chem- 
ist, in  an  endeavor  to  classify  the  properties  of  elements  and  their 
compounds,  arranged  the  former  in  the  order  of  their  atomic 
weights,  and  examined  them  from  this  point  of  view.  He  found 
that  starting  with  the  element  with  the  smallest  atomic  weight, 
hydrogen,  the  chemical  properties  of  the  succeeding  elements 
changed  markedly  as  the  atomic  weight  increased.  Lithium, 
which  has  the  atomic  weight  6.9,  is  an  active  base-forming  ele- 
ment which  has  the  valence  1 ;  the  next  element,  beryllium,  atomic 
weight  9.1,  has  less  active  base-forming  properties  and  has  the 
valence  2;  boron,  atomic  weight  11,  forms  weak  acids  and  has  the 
valence  3;  next  came  carbon,  atomic  weight  12,  which  is  also 
an  acid-forming  element  with  the  valence  4;  then  nitrogen,  14, 
with  the  valence  5,  oxygen,  16,  and,  finally,  fluorine,  19.  The 
next  element  which  was  sodium,  23,  resembled  lithium  markedly 
in  chemical  properties  and  had  the  same  valence  as  the  latter. 
Magnesium,  which  followed,  was  like  beryllium,  aluminium  like 
boron,  silicon  like  carbon,  phosphorus  like  nitrogen,  sulphur  like 
oxygen,  and  chlorine  like  fluorine.  Potassium,  the  element 
which  came  next,  was  the  first  member  of  a  series  of  seven  ele- 
ments the  properties  of  which  varied  with  increasing  atomic  weight 
in  the  same  way  as  the  properties  of  the  seven  elements  in  the 
series  from  lithium  to  fluorine,  inclusive,  and  those  in  the  series  of 
seven  elements  from  sodium  to  chlorine,  inclusive.  These  facts 


360 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


become  clearer  if  the  symbols  of  the  elements,  exclusive  of  hydrogen, 
and  their  atomic  weights  are  set  down  as  follows : 


Li,       6.9 

Be,       9.1 

B,     11. 

C,     12. 

N,     14. 

O,     16. 

F,     19. 

Na,    23. 

Mg,   24.3 

Al,   27.1 

Si,    28.3 

P,     31. 

S,     32. 

Cl,    35.5 

K,      39.1 

Ca,     40.1 

Sc,   44.1 

Ti,   48.1 

V,     51. 

Cr,   52. 

Mn,  54.9 

The  element  which  follows  manganese,  the  last  element  in  the  above 
tabulation,  is  iron,  55.9;  it  does  not  resemble  lithium,  sodium,  and 
potassium.  Nickel,  58.7,  and  cobalt,  59,  are  very  much  like  iron 
in  chemical  and  physical  properties.  Consequently,  these  three 
elements  were  set  off  by  themselves  and  became  members  of  the 
eighth  group  of  elements,  the  groups  consisting  of  the  members  in 
the  vertical  columns  in  the  above  tabulation;  they  are  numbered 
1  to  7  from  lithium  to  fluorine.  Copper,  63.6,  which  followed 
cobalt,  became  the  first  member  of  a  new  series  of  elements  the 
seventh  member  of  which  was  bromine,  and,  consequently,  fell 
into  its  place  under  chlorine  in  the  tabulation.  The  arrangement 
following  this  method  was  continued  in  the  same  way  with  the 
rest  of  the  elements,  and  led  to  a  table  which  resembled  closely 
that  given  facing  the  inside  of  the  back  cover  of  this  book.  Men- 
delejeff  put  forward  his  classification  of  the  elements  in  1869  and 
did  not  include  what  is  called  group  O  in  the  table,  as  these  elements, 
the  so-called  noble  gases,  were  unknown  at  that  time. 

From  what  has  been  brought  out  in  the  above  discussion  it  is 
evident  that  there  is  a  striking  relationship  between  the  physical 
and  chemical  properties  of  the  elements  and  their  atomic  weights, 
or  to  use  a  mathematical  term,  the  properties  of  the  elements  are  a 
function  of  their  atomic  weights.  They  are  not,  however,  a  direct 
function  of  the  atomic  weights,  for  as  the  latter  increase  there  is  a 
recurrence  in  properties  over  and  over  again;  for  this  reason  they 
are  said  to  be  a  periodic  function.  Mendelejeff's  study  of  the 
subject  led  him  to  this  conclusion,  and  when  he  published  his  first 
table  of  atomic  weights  he  proposed  the  periodic  law,  which  states 
that  the  physical  and  chemical  properties  of  the  elements  are  a 
periodic  function  of  their  atomic  weights. 

396.  A  few  additional  words  will  be  necessary  to  complete  the 
description  of  the  table  of  atomic  weights.  In  the  first  horizontal 


THE  PERIODIC  LAW  361 

line  under  the  numbers  of  the  groups  are  given  the  formulas 
E2O,  EO,  etc.  These  are  general  formulas  to  represent  the 
composition  of  the  oxides  of  the  members  of  the  group ;  for  example, 
in  group  3  we  have  boron,  scandium,  etc.,  the  oxides  of  which  have 
the  general  formula  E2O3,  that  is,  B2O3,  Sc2O3,  etc.  In  the 
second  horizontal  line,  we  have  beginning  with  group  4  the  symbols 
EH4,  EH3,  etc.  These  are  general  formulas  for  the  hydrides  of 
the  elements  in  the  several  groups,  such  as  CH4,  NH3,  etc.  In 
the  third  line  appear  the  letters  A  and  B  in  each  group.  These 
serve  to  designate  the  sub-groups  or  families  into  which  the  mem- 
bers of  the  groups  are  divided.  For  example,  while  there  is  a 
general  likeness  in  certain  respects  between  all  the  members  of 
group  2,  the  most  striking  similarities  appear  when  we  consider 
calcium,  strontium,  and  barium  on  the  one  hand,  and  magnesium, 
zinc,  cadmium,  and  mercury  on  the  other.  To  emphasize 
this  fact  the  symbols  of  one  set  of  elements  are  placed  on  one 
side  of  the  column  in  the  tabulation  and  those  of  the  other  on 
the  other  side.  The  dots,  .  .  .  .  ,  placed  in  some  of  the  squares 
indicate  the  fact  that  elements  to  fit  into  these  places  have  not  yet 
been  discovered.  When  Mendelejeff  proposed  this  classification 
of  the  elements,  the  elements  now  known  as  scandium,  group  3 
series  3,  gallium,  group  3  series  4,  and  germanium,  group  4  series  4, 
were  unknown.  If  the  tabulation  had  been  completed  according 
to  the  method  being  used,  when  titanium,  the  next  known  element 
after  calcium,  was  reached,  the  regularities  observed  in  the  first 
part  of  the  table  would  have  disappeared.  Mendelejeff  saw  that  if 
the  place  under  boron,  which  is  now  occupied  by  scandium,  was 
left  vacant,  titanium  would  fit  into  the  place  under  carbon  and  the 
regularities  would  not  be  disturbed.  He  did  this  and  prophesied 
that  an  element  with  an  atomic  weight  of  approximately  44  would 
be  discovered,  and  that  in  chemical  properties  and  in  the  com- 
position of  the  compounds  prepared  from  it,  it  would  resemble 
boron.  This  discovery  was  subsequently  made.  Mendelejeff 
left  vacant  the  spaces  now  occupied  by  gallium  and  germanium  for 
the  same  reasons,  and  his  prophecies  in  these  two  cases  were  also 
fulfilled.  These  facts  confirmed  in  a  striking  way  the  truth  under- 
lying the  periodic  classification  and  had  a  marked  influence  in 
leading  chemists  to  recognize  its  value. 

397.  When  the  system  was  proposed  the  atomic  weight  accepted 


362  INORGANIC  CHEMISTRY  FOR  COLLEGES 

for  uranium,  group  6  series  10,  was  119.2 — one-half  the  number 
assigned  to  it  to-day.  This  atomic  weight  would  place  the  ele- 
ment in  group  5  series  6  under  arsenic.  As  it  was  evident  that  its 
properties  were  not  in  accord  with  this  position,  Mendelejeff 
maintained  that  its  atomic  weight  should  be  doubled ;  he  did  this 
and  uranium  fell  into  the  place  under  molybdenum  and  tungsten, 
W,  where  it  belonged.  Subsequent  investigation  confirmed  the 
correctness  of  Mendelejeff 's  view.  It  was  evident  that  the 
atomic  weights  accepted  for  some  of  the  other  elements  were 
incorrect;  these  were  placed  in  the  table  in  accordance  with  their 
properties,  and  in  nearly  every  case  redetermination  of  the  values 
led  to  a  confirmation  of  the  correctness  of  the  conclusions  indicated 
by  the  periodic  law.  In  the  case  of  tellurium,  group  6  series  6, 
however,  this  has  not  been  the  case.  It  belongs,  evidently,  to 
the  sulphur-selenium  family  and  it  is  placed  in  the  position  to  indi- 
cate this  fact,  although  its  atomic  weight  being  greater  than  iodine, 
puts  it  in  group  8. 

The  periodic  law  has  been  a  guide  of  the  greatest  value  in 
chemical  investigation  for  over  fifty  years;  it  has  suggested  the 
possibility  of  the  preparation  of  many  compounds  with  valuable 
properties,  and,  after  their  discovery,  such  compounds  have  been 
studied  from  the  standpoint  of  the  law.  Its  greatest  value,  how- 
ever, is,  perhaps,  the  service  rendered  in  assisting  in  the  classifica- 
tion of  the  many  facts  of  inorganic  chemistry.  The  student  will 
find  it  helpful  in  remembering  many  of  these  facts.  Since  science 
is  systematized  or  classified  knowledge,  the  periodic  law  has  been 
a  most  important  factor  in  the  development  of  the  science  of 
chemistry. 

As  the  elements  and  their  compounds  have  been  intensively 
studied  in  recent  years,  many  facts  have  been  discovered  which 
are  in  striking  accord  with  the  periodic  classification;  but  others  in 
an  equally  striking  manner  cannot  be  interpreted  through  the  use 
of  this  generalization.  These  exceptions  do  not  lead  the  chemist 
to  refuse  to  accept  the  law  as  an  expression  of  a  truth,  but  rather 
inspire  him  to  a  more  detailed  study  of  the  facts  in  the  hope  of 
discovering  the  causes  of  the  exceptions,  in  order  to  arrive  at  a 
more  perfect  expression  of  the  truth  in  a  modified  law. 

398.  The  periodic  law  considers  the  elements  solely  from  the 
standpoint  of  the  weight  of  the  atoms  of  the  elements.  It  is 


THE  PERIODIC  LAW  363 

evident  that  it  must  be  an  imperfect  expression  of  the  truth  in 
regard  to  the  properties  of  the  elements  if  their  content  of  energy 
is  a  factor  in  determining  their  properties.  The  study  of  the  atom 
from  this  point  of  view  began  with  the  discovery  of  radium.  A 
great  deal  of  evidence  has  been  found  for  the  hypothesis  that  the 
atoms  of  the  elements  are  not  single  units  but  are  made  up  of 
so-called  electrons  which  are  charges  of  negative  electricity  that 
surround  a  nucleus  having  a  positive  charge.  Since  the  atom 
contains  positive  and  negative  electricity  it  is  a  storehouse  of 
energy ;  and  when  it  unites  with  another  atom  a  part  of  this  energy 
is  transformed  in  the  chemical  union.  According  to  this  view 
the  elements  have  been  built  up  by  progressively  increasing  the 
number  of  positive  charges  and  negative  charges  which  go  to  make 
up  the  atoms  of  the  several  elements.  The  periodic  classification 
of  the  elements  which  results  from  the  conception  of  the  electrical 
constitution  of  the  atom  is  more  in  accord  with  the  facts  than  that 
arrived  at  from  the  consideration  of  the  weight  relations;  the 
exceptions  to  the  law  of  Mendelejeff  disappear,  and  many  facts 
which  cannot  be  interpreted  by  the  law  or  could  not  be  foretold 
by  it  are  in  accordance  with  the  newer  classification  (798) . 

The  periodic  classification  of  the  elements  cannot  be  advan- 
tageously discussed  in  detail  without  a  knowledge  of  a  large  number 
of  facts.  For  this  reason  it  will  be  considered  as  the  chemistry 
of  the  elements  and  their  compounds  is  developed  in  subsequent 
chapters.  Its  application  to  the  facts  as  they  are  brought  forward 
will  bring  out  clearly  the  importance  of  the  law  as  an  aid  in  their 
generalization.  The  classification  according  to  Mendelejeff  will 
be  used,  as  it  serves  adequately  in  the  consideration  of  the  facts 
to  be  presented. 


CHAPTER  XXVI 
THE  HALOGEN  FAMILY 

399.  The  members  of  the  so-called  halogen  family  constitute 
sub-group  B  in  the  seventh  group  in  the  periodic  classification  of 
the  elements.  A  study  of  the  members  of  this  sub-group  from 
the  point  of  view  of  searching  out  similarities  and  differences  in 
chemical  and  physical  properties,  will  bring  out  clearly  the  sig- 
nificance of  the  periodic  law ;  and  the  student  will  see  more  clearly 
than  in  the  past  how  the  many  facts  of  chemistry  may  be  corre- 
lated and  thus  become  examples  of  general  principles.  A  compari- 
son of  the  chemistry  of  bromine  with  the  chemistry  of  chlorine 
will  necessitate  a  review  of  the  facts  concerning  the  latter,  which 
will  thus  be  impressed  on  the  memory  more  definitely.  Repe- 
tition is  the  most  important  factor  in  memory,  and  when  the  repe- 
tition of  a  set  of  facts  is  associated  with  the  acquisition  of  new  facts 
of  importance  and  with  the  discovery  of  relationships  of  interest, 
the  process  becomes  a  source  of  mental  satisfaction.  The  student 
should  strive  to  study  in  this  way,  and  when  the  method  has 
become  a  habit  of  thought  as  the  result  of  repetition,  one  of  the 
attributes  of  a  trained  mind  will  have  been  attained,  and  one  of 
the  objects  of  education  accomplished. 

The  members  of  the  halogen  family  are  active  acid-forming, 
electro-negative  elements.  Their  activity  in  this  respect  decreases 
with  increasing  atomic  weight.  This  is  evident  from  a  consider- 
ation of  the  change  in  energy  which  takes  place  when  the  elements 
react  with  hydrogen,  a  typical  electro-positive  element.  The 
heats  of  formation  (165)  of  the  halogen  hydrides,  HF,  HC1,  HBr, 
HI  are,  respectively,  38.5,  22,  8.5,  and  0.5  (at  400°)  large  calories. 
The  values  are  those  obtained  in  the  production  of  1  gram-molec- 
ular-weight of  the  hydrides  from  hydrogen  and  the  halogens  in 
the  gaseous  condition.  These  values  indicate  clearly  the  order  in 
which  the  elements  stand  when  arranged  according  to  their 

364 


THE  HALOGEN  FAMILY  365 

chemical  activity.  This  order  is  just  the  reverse  of  that  arrived 
at  if  the  arrangement  is  made  according  to  the  weight  relations. 
This  means  that  the  activity  of  the  elements  decreases  as  their 
atomic  weights  increase.  This  relationship  is  shown  in  the  case 
of  most  families  of  elements — the  lighter  atoms  are  the  more 
active.  We  shall  see  as  the  subject  develops,  that  if  we  arrange 
these  elements  or  their  compounds  in  which  the  former  play  the 
role  of  electro-negative  elements,  according  to  any  property  which 
results  from  chemical  activity  or  from  the  weight  of  the  atoms, 
the  order  will,  in  most  cases,  be  the  same,  namely,  fluorine,  chlorine, 
bromine,  and  iodine.  The  facts  about  to  be  presented  should  be 
examined  from  this  point  of  view, 

BROMINE 

400.  Occurrence  and  Discovery. — Bromine  is  found  chiefly 
as  sodium  bromide  and  magnesium  bromide  associated  with  the 
chlorides  of  these  elements.  As  the  salts  are  soluble  they  find 
their  way  into  the  ocean,  and  the  chief  sources  of  the  halogen  are 
salt  mines  that  have,  in  all  probability,  been  formed  as  the  result 
of  the  drying  up  of  inland  seas.  The  quantity  of  bromides  found 
is  very  small  compared  with  that  of  the  chlorides. 

The  element  was  discovered  by  Balard  in  1826  as  the  result 
of  treating  with  chlorine  a  solution  from  which  crude  salt  had  been 
crystallized — the  so-called  mother-liquor.  The  bromides  present 
in  the  latter  reacted  with  chlorine  and  formed  chlorides,  and 
bromine  was  set  free.  The  reaction  in  the  case  of  sodium  bromide 
is  as  follows: 

2NaBr  +  C12  =  2NaCl  +  Br2 

A  deep  yellow  color  was  formed  in  the  solution.  Balard  shook  the 
latter  with  ether,  which  dissolved  the  colored  material  and  thus 
separated  it  from  the  water.  On  evaporation  of  the  ether  the 
residue  left  was  found  to  be  a  red  liquid.  On  account  of  its 
marked  odor  the  name  selected  for  it  was  derived  from  the  Greek 
word  signifying  a  stench. 

Salt  beds  which  contain  bromine  in  the  proportion  which  makes 
it  profitable  to  use  them  as  a  source  of  bromine  occur  at  Stassfurt 
in  Germany,  and  in  the  States  of  Michigan,  Ohio,  and  West 
Virginia. 


366  INORGANIC  CHEMISTRY  FOR  COLLEGES 

401.  Preparation  of  Bromine.  —  In  the  manufacture  of  bromine 
on  the  industrial  scale  the  reaction  discovered  by  Balard  is  gen- 
erally used.  The  brine,  obtained  by  dissolving  the  crude  salt  in 
water,  is  treated  with  just  enough  chlorine  to  liberate  the  bromine 
present.  Steam  is  then  passed  through  the  liquid  and  the  bro- 
mine passes  off  and  is  condensed  along  with  the  steam,  and  sep- 
arates as  a  heavy,  reddish-brown  liquid. 

Bromine  can  be  separated  from  its  compounds  by  utilizing  the 
methods  which  have  been  described  in  connection  with  the  prep- 
aration of  chlorine.  It  is  formed,  for  example,  when  manganese 
dioxide  is  treated  with  hydrobromic  acid, 

MnO2  +  4HBr  =  MnBr2  +  2H2O  +  Br2 

and  when  a  bromide  is  gently  heated  with  manganese  dioxide  and 
concentrated  sulphuric  acid: 

2NaBr  +  2H2SO4  +  MnO2  =  Na2SO4  +  MnSO4  +  2H2O  +  Br2 

The  electrolysis  of  bromides  also  yields  bromine  just  as  that  of 
chlorides  yields  chlorine. 

The  reaction  by  which  a  bromide  is  converted  into  a  chloride 
when  treated  with  free  chlorine  is  an  important  one.  The  fact 
that  the  reaction  takes  place  shows  clearly  that  chlorine  is  a  more 
active  element  than  bromine.  When  the  reaction  takes  place  in 
solution  it  occurs  between  bromine  ions  and  chlorine  molecules 
and  indicates  that  the  latter  has  a  greater  tendency  to  exist  as  an 
ion  than  the  former  has.  The  equation  for  the  reaction, 

2Na+  +  2Br~  +  C12  =  2Na+  +  Br2  +  2CP 

resembles  in  this  respect  that  of  the  reaction  between  salts  and 
metals,  for  example,  the  action  of  iron  on  a  solution  of  a  copper  salt  : 

Cu+  f  +  2Cr  +  Fe  =  Cu  +  2C1~  +  Fe+  + 


In  the  latter  case  we  have  learned  that  the  reaction  is  said  to  take 
place  because  iron  has  a  greater  tendency  than  has  copper  to  be 
an  ion  —  its  solution  pressure  is  higher.  We  now  see  that  elements 
that  pass  into  solution  as  negative  ions  behave  in  an  analogous 
manner.  Chlorine  has  a  greater  solution  pressure  than  bromine. 
The  elements  that  form  negative  ions  can  be  arranged  in  a  series 


THE  HALOGEN  FAMILY  367 

based  on  this  property  in  the  same  way  as  the  metallic  elements 
(574). 

402.  Physical  Properties. — Bromine,  atomic  weight  79.92,  is  a 
reddish-brown  liquid  (sp.  gr.  3.12  at  20°),  which  boils  at  59°,  and 
freezes  to  needle-shaped  crystals  at   —7.3°;    it  possesses  a  most 
disagreeable  and  pungent  odor.     One  hundred  cubic  centimeters 
of  water  dissolve  4.3  grams  of  the  liquid  at  0°,  and  about  3  grams 
at  ordinary   temperatures;    the  solution,   called  bromine-water, 
gives  off  the  halogen  freely  in  the  form  of  a  vapor,  which  has  a 
reddish  color.       Bromine  has  a  high  vapor  pressure  at  room- 
temperature;    at  18°  the  pressure  is  150  mm.,  and,  as  a  conse- 
quence, the  gas  over  liquid  bromine  contained  in  a  bottle  at  this 
temperature  is  approximately  one-fifth  bromine  and  four-fifths  air. 

403.  Chemical  Properties. — Bromine  affects  markedly  the  tis- 
sues of  which  the  body  is  made  up.     It  has  a  very  irritating  effect 
on  mucous  membrane,  and  the  inhalation  of  its  vapor  produces 
most  disagreeable  effects.     When  left  in  contact  with  the  skin 
bromine  causes  painful  sores  which  heal  slowly. 

In  chemical  conduct  bromine  closely  resembles  chlorine;  it 
combines  directly  with  metals  and  many  non-metals  to  form 
bromides,  the  properties  of  which  are  similar  to  those  of  the  cor- 
responding chlorides.  It  forms  at  low  temperatures  a  hydrate 
with  water,  Br2,10H2O,  which  is  less  stable  than  the  hydrate  of 
chlorine. 

404.  Uses. — Bromine  could  be  used  for  many  of  the  purposes 
to  which  chlorine  is  put,  but  this  is  not  done  unless  there  are  some 
particular    advantages  which    come    from  its  use.     Bromine   is 
employed  in  the  preparation  of  bromides,  some  of  which  are  exten- 
sively used.     Potassium  bromide  is  used  in  medicine  as  a  sedative, 
and  silver  bromide  is  the  most  important  substance  in  the  sensitive 
film  of  a  photographic  plate. 

Free  bromine  is  used  in  certain  metallurgical  processes  and  in 
the  preparation  of  several  important  dyes  and  other  organic 
substances. 

HYDROBROMIC  ACID 

•  405.  Preparation. — Hydrobromic  acid,  which  is  a  colorless  gas, 
is  not  made  in  a  way  analogous  to  that  used  for  preparing  hydro- 
chloric acid,  namely,  by  treating  a  bromide  with  concentrated 


368  INORGANIC  CHEMISTRY  FOR  COLLEGES 

sulphuric  acid.  A  similar  reaction  takes  place  in  this  case,  but 
the  acid  formed  is  in  part  oxidized  by  sulphuric  acid  to  bromine 
and  water.  It  can  be  prepared  conveniently  by  reactions  entirely 
analogous  to  those  into  which  chlorine  enters,  but  which  are  not 
used  for  the  preparation  of  hydrochloric  acid.  Bromine  reacts 
with  hydrocarbons  and  forms  substitution-products  and  the  acid. 
Benzene,  CeHe,  can  be  used  for  this  purpose : 

C6H6  +  Br2  =  C6H5Br  +  HBr 

A  similar  reaction  takes  place  smoothly  in  the  cold  when  anthra- 
cene, CnHio,  a  product  derived  from  coal-tar,  is  used. 

Hydrobromic  acid  is  conveniently  prepared  by  the  action  of 
water  on  the  bromides  of  phosphorus.  When  the  acid  is  made  in 
this  way  bromine  is  first  allowed  to  drop  slowly  on  red  phosphorus, 
which  is  converted  into  phosphorus  pentabromide;  water  is  next 
added,  drop  by  drop,  and  as  a  result  hydrobromic  acid  is  formed: 

PBr3  +  3H2O  =  HsP03  +  3HBr 

406.  Properties. — Hydrobromic  acid  is  a  colorless  gas,  which 
fumes  in  the  air,  and  has  a  sharp  odor.  By  reduction  in  tempera- 
ture it  can  be  converted  into  a  liquid,  which  boils  at  —69°.  The 
gas  is  very  soluble  in  water,  600  volumes  dissolving  in  1  volume  of 
the  latter  at  0°. 

The  methods  described  above  are  used  when  hydrobromic  acid 
is  required  in  the  form  of  a  gas;  if  an  aqueous  solution  is  desired 
the  reaction  between  a  bromide  and  sulphuric  acid  may  be  used. 
This  is  possible  because  the  oxidizing  power  of  sulphuric  acid  is 
markedly  reduced  in  the  presence  of  water.  It  will  be  recalled 
that  this  fact  is  an  example  of  a  general  truth;  ions,  in  general, 
do  not  act  as  oxidizing  agents,  and  when  sulphuric  acid  is  dissolved 
in  water  the  molecules  are  largely  ionized.  When  it  is  desired 
to  make  an  aqueous  solution  of  hydrobromic  acid,  potassium  or 
sodium  bromide,  water,  and  sulphuric  acid,  in  the  correct  propor- 
tions, are  brought  together  and  the  mixture  is  distilled.  The 
acid  and  water  form  a  constant-boiling  mixture  (142)  which  con- 
tains approximately  48  per  cent  of  HBr,  has  the  specific  gravity 
1.49,  and  boils  at  126°  at  760  mm.  pressure.  In  making  the  acid 
in  this  way  the  materials  used  are  mixed  in  the  proportion  of  1 
mol  of  the  salt,  1  of  sulphuric  acid,  and  5  of  water,  and  distilled. 


THE  HALOGEN  FAMILY  369 

407.  It  has  been  pointed  out  that  bromine  is  a  less  active  ele- 
ment than  chlorine;  its  compound  with  hydrogen  is,  thus,  more 
readily  decomposed  than  hydrochloric  acid.  This  is  clearly  seen  in 
its  behavior  with  oxidizing  agents.  Concentrated  sulphuric  acid, 
which  is  not  sufficiently  active  as  an  oxidizing  agent  to  affect 
hydrochloric  acid,  oxidizes  hydrobromic  acid.  The  equation  for 
the  reaction  may  be  analyzed  as  follows: 

H2S04  =  H20  +  SO2  +  [O] 
2HBr  +  [O]  =  H2O  +  Br2 


H2SO4  +  2HBr  «=>  2H2O  +  SO2  +  Br2 

408.  Properties  of  Bromides. — The  bromides  of  the  common 
metallic  elements  resemble  in  physical  properties  the  corresponding 
chlorides.     They  are  all  more  or  less  easily  soluble  in  water  except 
those  of  silver,  lead,  and  mercury  (mercurous  bromide,  HgBr). 
They  are  less  stable  towards  heat  than  the  chlorides. 

The  bromides  of  the  acid-forming  elements  also  resemble 
the  chlorides  of  these  elements.  They  have  higher  boiling-points, 
and  are  decomposed  by  water  with  the  formation  of  hydrobromic 
acid  and  the  acid  derived  from  the  other  electro-negative  element 
which  they  contain. 

409.  Test  for  Bromides. — When  a  solution  of  silver  nitrate  is 
added  to  a  solution  of  a  bromide,  silver  bromide  is  precipitated  in 
the  form  of  a  cloud  or  as  a  curdy  solid.     The  solid,  which  has  a  light 
yellow   tint,    is   insoluble   in   nitric   acid.     The   reaction,    which 
resembles  closely  that  which  takes  place  in  the  case  of  a  chloride, 
is  as  follows : 

NaBr  +  AgNO3  =  AgBr  +  NaNO3 

Silver  chloride  is  readily  soluble  in  dilute  ammonia,  whereas  silver 
bromide  is  difficultly  soluble. 

If  an  unknown  substance  behaves  with  silver  nitrate  in  a  way 
to  indicate  that  it  is  a  bromide,  another  portion  of  its  solution  is 
treated,  drop  by  drop,  with  chlorine-water,  and  shaken  with  2  or  3 
c.c.  of  carbon  disulphide  or  chloroform.  The  bromine  set  free  in 
the  reaction, 

2NaBr  +  C12  =  2NaCl  +  Br2, 


370  INORGANIC  CHEMISTRY  FOR  COLLEGES 

passes  into  the  carbon  disulphide  and  becomes  evident  on  account 
of  its  red  color.  This  behavior  is  shown  by  all  substances  which 
yield  bromine  ions  when  dissolved  in  water.  In  adding  chlorine- 
water  to  a  solution  of  a  bromide  in  making  the  test,  it  is  necessary 
to  avoid  the  use  of  more  chlorine  than  is  required  to  liberate  the 
bromine — an  excess  of  the  reagent;  in  the  presence  of  water, 
chlorine  converts  bromine  into  a  colorless  compound  (431). 

IODINE 

410.  Occurrence  and  Discovery. — Iodine  occurs  in  the  com- 
bined state  in  sea-water  in  very  small  quantities.     It  is  present 
in  certain  aquatic  plants  and  animals;   and  the  former  have  been 
for  many  years  a  source  of  the  element.     In  1811  Courtois,  a 
French  manufacturer  of  soap,  carried  out  some  experiments  with 
one  of  the  products  used  in  making  soap,  namely,  the  ashes  obtained 
by  burning  kelp,  which  is  a  variety  of  sea-weed,  that  was  used 
as  a  source  of  the  potash  required.     He  treated  the  ashes  with 
concentrated  sulphuric  acid  and  discovered,  as  a  result,  that  a 
violet-colored  vapor  was  given  off.     Later,  in  1814.  Gay-Lussac 
published  a  full  account  of  the  substance  discovered  by  Courtois, 
and  proved  that  it  was  an  element.     The  name  selected  for  the 
newly  discovered  element,  iodine,  was  derived  from  the  Greek 
word  meaning  violet. 

Iodine  is  still  obtained  from  the  ashes  of  kelp,  which  contain 
from  0.5  to  1.5  per  cent  of  a  mixture  of  sodium  and  potassium 
iodides.  The  chief  source  at  present  is  crude  Chile  saltpeter, 
NaNO3,  which  contains  about  0.2  per  cent  of  sodium  iodate, 
NaIO3. 

Iodine  occurs  in  the  human  body  in  traces,  its  presence  in  the 
thyroid  gland  being  of  great  significance.  In  the  disease  known 
as  goiter  this  gland  is  ill  developed;  medical  research  has  shown 
that  if  a  substance,  called  iodothyrine,  which  is  obtained  from  the 
thyroid  of  the  sheep,  is  injected,  the  human  gland  develops  and  the 
disease  which  accompanies  its  inactivity  disappears  to  a  large 
extent. 

411.  Preparation. — Iodine  is  prepared  from  iodides  by  reac- 
tions entirely  analogous  to  those  used  in  the  preparation  of  bromine. 
In  the  industrial  preparation  from  the  ashes  of  kelp,  the  latter  are 


THE  HALOGEN  FAMILY  371 

treated  with  manganese  dioxide  and  sulphuric  acid.  The  mixture 
is  heated  to  the  temperature  at  which  the  halogen  sublimes,  and  the 
violet-colored  vapor  formed  condenses  to  black  crystals.  In  the 
laboratory  potassium  iodide  is  usually  used  in  preparing  iodine  in 
this  way.  The  iodine  can  also  be  liberated  by  treating  the  iodide 
with  chlorine,  and  this  process  is  used  industrially.  In  the  man- 
ufacture of  iodine  from  crude  Chile  saltpeter  the  material  is  dis- 
solved in  water  and  treated  with  a  mixture  of  sodium  sulphite  and 
sodium  bisulphite.  To  obtain  iodine  from  sodium  iodate,  NalOa, 
it  is  necessary  to  use  a  reducing  agent  to  remove  the  oxygen  from 
the  compound;  the  sulphites  serve  this  purpose.  The  iodine 
which  precipitates  as  a  black  solid,  is  filtered  off,  dried,  and 
sublimed. 

In  purifying  iodine  it  is  mixed  with  potassium  iodide  and 
resublimed.  If  any  chlorine  is  present  in  the  iodine,  as  iodine 
chloride,  IC1,  the  latter  reacts  with  the  iodide  as  follows : 

KI  +  IC1  =  KC1  +  I2 

412.  Physical  Properties. — Iodine,  atomic  weight  126.92,  is 
obtained  in  the  form  of  black  crystalline  plates  when  its  vapor  is 
condensed.  It  has  the  specific  gravity  4.93,  melts  at  114°,  and 
boils  at  184°.  The  vapor  of  iodine  has  a  violet  color  and  is  irri- 
tating when  inhaled.  A  saturated  aqueous  solution  of  iodine 
contains  0.32  gram  of  the  halogen  per  liter  and  has  a  light  brown 
color;  it  is  more  soluble  in  alcohol,  ether,  chloroform,  and  carbon 
disulphide.  Solutions  of  iodine  in  solvents  which  contain  oxygen, 
such  as  water,  ether,  and  alcohol,  are  brown  in  color,  whereas 
solutions  in  other  solvents  are  violet.  The  brown  color  is  thought 
to  be  due  to  the  presence  in  the  first  class  of  solvents  of  more 
or  less  stable  addition-products  of  the  solvent  and  the  dissolved 
iodine.  In  the  case  of  solvents  which  contain  no  oxygen  the  halo- 
gen has  the  same  color  that  it  would  have  if  it  were  in  the  form  of 
vapor. 

Iodine  is  very  soluble  in  solutions  of  potassium  iodide  and 
other  iodides,  the  amount  dissolving  being  determined  by  the 
concentration  of  the  iodide  present.  Such  solutions  are  brown  in 
color,  and  it  has  been  shown  that  a  compound  is  present  which  is 
formed  as  the  result  of  the  direct  addition  of  the  iodide  and  iodine: 

KI  +  I2  <=»  KI3 


372  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  reaction  is  a  reversible  one,  for  if  a  solution  of  potassium 
iodide  is  saturated  with  iodine,  and  then  diluted,  a  part  of  the 
latter  is  precipitated.  A  solution  of  iodine  in  potassium  iodide  is 
used  in  many  processes  in  analytical  chemistry.  The  indicator 
used  in  such  processes,  which  are  classed  together  under  the  name 
iodimetry,  is  a  dilute  solution  of  starch.  When  a  few  cubic 
centimeters  of  such  a  solution  are  added  to  a  solution  containing 
free  iodine,  a  characteristic  blue  color  is  produced.  It  is  probable 
that  the  color  is  formed  as  the  result  of  the  physical  adsorption  of 
the  halogen  by  the  starch.  The  reaction  which  serves  to  show  the 
presence  of  exceedingly  small  amounts  of  iodine  is  also  used  to  test 
for  the  presence  of  starch;  in  this  case  a  dilute  solution  of  iodine 
is  added  to  the  material  tested. 

413.  Chemical  Properties. — The  family  relationship  between 
chlorine,  bromine,  and  iodine  is  clearly  shown  in  the  chemical 
properties  of  the  latter.  It  is  much  less  active  than  the  other 
halogens,  and  any  differences  in  behavior  can  be  traced  to  this 
cause.  Iodine  forms  iodides  with  metals  and  certain  non-metals, 
but  these  are  less  stable  than  the  corresponding  bromides.  It 
combines  with  hydrogen  very  slowly  at  elevated  temperatures 
to  form  hydrogen  iodide.  The  affinity  of  iodine  for  hydrogen  is  so 
small  that  under  ordinary  conditions  it  does  not  react  with  hydro- 
carbons in  the  way  that  chlorine  and  bromine  do. 

The  water-solution  of  iodine  shows  some  of  the  properties  of 
similar  solutions  of  bromine  and  chlorine,  but  the  proportion  of 
hypoiodous  acid  formed  is  excessively  small,  and  its  presence 
becomes  evident  only  when  something  is  brought  into  contact 
with  it  with  which  it  will  react.  For  example,  a  solution  of  sul- 
phur dioxide  in  water  (sulphurous  acid)  will  decolorize  a  solution 
of  iodine: 

I2  +  H2O  <=±  [HIO]  +  HI 
[HIO]  +  H2SO3  <=*  HI  +  H2SO4 

I2  +  H2O  +  H2SO3  <=*  2HI  +  H2SO4 

The  reaction  is  a  reversible  one  and  if  it  is  to  run  to  completion 
from  left  to  right,  the  solution  must  be  dilute  and  an  excess  of 
iodine  used.  The  reaction  is  commonly  employed  in  determining 
quantitatively  the  amount  of  sulphurous  acid  present  in  a  given 


THE  HALOGEN  FAMILY  373 

solution.  The  amount  of  iodine  added  can  be  calculated  if  a 
definite  volume  of  a  solution  of  iodine  of  known  strength  is  used; 
and  the  amount  left  over  can  be  determined  by  titrating  the  solu- 
tion with  a  solution  of  sodium  thiosulphate.  The  reaction  which 
takes  place  is  as  follows : 

I2  +  2Na2S2O3  =  2NaI  +  Na2S406 

The  difference  between  the  amounts  of  iodine  used  in  the  two  opera- 
tions is  the  quantity  which  was  required  to  oxidize  the  sulphurous 
acid,  and  is,  therefore,  a  measure  of  the  amount  of  the  latter. 
Many  substances  are  oxidized  in  aqueous  solution  by  iodine,  and 
the  process  outlined  above  is  an  important  one  in  quantitative 
analysis. 

Iodine  in  the  presence  of  water  reacts  with  other  substances 
that  are  easily  oxidized.  For  example,  hydrogen  sulphide  and 
iodine  form  hydriodic  acid  and  sulphur: 

H2S  +  I2  =  2HI  +  S 

The  reaction  expressed  by  the  above  equation  may  be  utilized 
to  prepare  an  aqueous  solution  of  hydriodic  acid.  A  solution  of 
hydrobromic  acid  can  be  made  by  an  analogous  reaction.  For 
this  purpose  the  halogen  is  covered  with  the  amount  of  water 
desired  as  the  solvent  for  the  acid  formed,  and  a  stream  of  hydrogen 
sulphide  gas  is  passed  into  the  solution.  When  the  reaction  is 
complete  the  sulphur  is  filtered  off  and  the  liquid  is  distilled. 

414.  Uses. — Iodine  affects  the  skin  and  mucous  membrane, 
but  its  action  is  not  so  marked  as  that  of  bromine.     It  reduces 
swellings  and  hardens  the  skin.     It  was  much  used  in  the  army  to 
relieve  the  pain  which  results  from  abrasions  of  the  feet  induced 
by  long  marches.     The  affected  parts  were  treated  with  an  alco- 
holic solution  (tincture)  of  iodine. 

Iodine,  as  well  as  a  number  of  organic  compounds  which  contain 
the  element,  has  marked  antiseptic  properties;  iodoform,  CHIs, 
which  is  extensively  used  in  surgery  for  this  reason,  is  manufac- 
tured from  alcohol,  sodium  hydroxide,  and  iodine.  Sodium, 
potassium,  and  rubidium  iodides  are  used  in  medicine. 

HYDRIODIC  ACID 

415.  Preparation. — Hydriodic  acid,  which  is  a  colorless  gas,  is 
not  prepared  by  the  action  of  sulphuric  acid  on  an  iodide;  even  in 


374  INORGANIC  CHEMISTRY  FOR  COLLEGES 

aqueous  solution  the  acid  oxidizes  the  hydriodic  acid  formed  and 
iodine  is  set  free.  It  cannot  be  prepared  by  the  action  of  iodine 
on  a  hydrocarbon  in  a  way  analogous  to  that  by  which  hydrogen 
bromide  can  be  made.  The  gas  can  be  prepared  by  heating 
sodium  iodide  with  concentrated  phosphoric  acid,  HaPCU.  In 
this  case  a  double  decomposition  takes  place;  as  an  excess  of  the 
acid  is  used  an  acid  phosphate  is  formed : 

Nal  +  H3PO4  =  NaH2PO4  +  HI 

We  shall  learn  later  that  phosphoric  acid  is  not  an  oxidizing  agent; 
as  a  consequence,  it  does  not  react  with  hydrogen  iodide. 

The  method  commonly  employed  to  make  the  gas  is  to  treat 
with  water,  phosphorus  iodide  prepared  by  the  action  of  iodine  on 
phosphorus : 

PI3  +  3H20  =  P(OH)3  +  SHI 

416.  Physical  Properties. — Hydriodic  acid  is  a  heavy,  colorless 
gas  which  fumes  in  the  air.     It  can  be  condensed  to  a  liquid  which 
boils  at  -34°  and  freezes  at  -51°.     One  volume  of  water  at  10° 
dissolves  425  volumes  of  the  gas,  and  the  resulting  solution  con- 
tains 70  per  cent  of  hydrogen  iodide.     It  forms  a  constant-boiling 
mixture  with  water,  which  has  the  specific  gravity  1.7,  contains 
57  per  cent  of  hydrogen  iodide,  and  boils  at  127°  at  760  mm.  pres- 
sure. 

417.  Chemical  Properties. — Hydriodic  acid  resembles  in  its 
general  behavior  hydrochloric  acid  and  hydrobromic  acid;    the 
differences  observed  are  due  to  the  fact  that  the  latter  are  much 
the  more  stable  substances.    Hydrogen  iodide  begins  to  decompose 
slowly  into  hydrogen  and  iodine  at  180°;  at  higher  temperatures 
the  reaction  is  a  rapid  one.     This  fact  can  be  easily  shown  by 
pouring  from  a  large  jar  a  quantity  of  the  gas  on  a  Bunsen  flame. 
As  the  heavy  gas  falls  on  the  flame  it  becomes  heated,  is  decom- 
posed, and  a  violet  cloud  of  iodine  vapor  is  seen. 

If  gaseous  hydriodic  acid  is  mixed  with  chlorine  a  reaction 
takes  place  with  the  evolution  of  light: 

2HI  +  C12  =  2HC1  +  I2 

In  aqueous  solution  both  the  acid  and  its  salts  are  decomposed  by 
chlorine  in  this  way;  bromine  acts  in  a  similar  manner. 


THE  HALOGEN  FAMILY  375 

Hydriodic  acid  is  a  strong  acid,  and  resembles  closely  hydro- 
chloric acid.  Its  solution  is  unstable  in  the  presence  of  the  air  on 
account  of  the  fact  that  oxygen  reacts  with  it  slowly  at  ordinary 
temperatures : 

4HI  +  O2  =  2H2O  +  2I2 

The  reaction  takes  place  more  rapidly  in  the  presence  of  sunlight, 
and  to  decrease  the  rate  of  the  decomposition  the  solution  is  usually 
kept  in  an  opaque  container  or  a  bottle  of  brown  glass.  Oxidizing 
agents  liberate  iodine  from  hydriodic  acid.  Its  behavior  with 
sulphuric  acid  has  been  noted.  When  solid  potassium  iodide  is 
treated  with  concentrated  sulphuric  acid,  the  presence  of  sulphur 
dioxide,  sulphur,  and  hydrogen  sulphide  can  be  noted ;  there  is  also 
present  some  hydrogen  iodide  which  escapes  decomposition. 

A  dilute  solution  of  hydriodic  acid  is  used  in  medicine  to  pro- 
mote certain  secretions.  The  acid  is  used  in  organic  chemistry  to 
prepare  iodides  and  to  test  for  the  presence  of  certain  important 
groupings  of  elements;  it  is  one  of  the  most  active  reducing 
agents  and  is  used  for  this  purpose  when  other  agents  fail. 

418.  Test  for  Iodides. — The  reactions  upon  which  the  test  for 
iodides  is  based  are  analogous  to  those  used  in  the  case  of  bromides. 
When  silver  nitrate  is  added  to  a  solution  of  an  iodide  a  precipitate 
of  silver  iodide  is  formed,  which  is  insoluble  in  nitric  acid  and 
ammonia,  and  has  a  greenish-yellow  color.    If  the  behavior  of  a  salt 
with  silver  nitrate  indicates  that  it  is  an  iodide,  to  a  second  portion 
of  the  solution  is  added  a  few  cubic  centimeters  of  carbon  disul- 
phide;  chlorine-water  is  next  added  drop  by  drop,  and  the  solution 
shaken.     If  an  iodide  is  present  iodine  is  set  free,  dissolves  in  the 
carbon  disulphide,  and  is  evident  as  the  result  of  the  violet  color 
of  its  solution. 

If  a  solution  contains  both  a  bromide  and  an  iodide,  the  iodine 
is  first  set  free  as  the  chlorine-water  is  slowly  added;  when 
this  has  reacted  with  the  added  chlorine  to  form  a  colorless  com- 
pound (431),  the  bromine  is  liberated  and  imparts  to  the  carbon 
disulphide  its  characteristic  red  color. 

FLUORINE 

419.  The  element  which  has  the  smallest  atomic  weight  in  the 
family  of  the  halogens  is  fluorine,  and  on  this  account  it  might 


376  INORGANIC  CHEMISTRY  FOR  COLLEGES 

have  been  considered  first  in  a  discussion  of  the  group.  Its  con- 
sideration has  been  delayed,  however,  for  a  good  reason.  Fluorine 
is  an  exceedingly  active  element  and,  consequently,  it  enters  into 
reactions  not  exhibited  by  chlorine,  and  its  compounds  are  so 
stable  that  many  of  them  do  not  show  the  behavior  exhibited  by 
the  chlorides  of  similar  composition.  Chlorine,  on  the  other  hand, 
as  we  have  seen,  does  not  differ  markedly  in  activity  from  bromine 
and  a  consideration  of  the  latter  after  the  chemistry  of  chlorine 
has  been  mastered  gives  an  excellent  opportunity  to  contrast  the 
behavior  of  the  two  elements  and  their  compounds,  and  to  bring 
out  relationships  indicated  by  the  periodic  law.  We  have  also 
seen  that  while  iodine  resembles  chlorine  and  bromine  in  many 
respects,  its  lack  of  activity  results  in  a  behavior  which  is  char- 
acteristic. We  are  now  in  a  position  to  study  the  chemistry  of  an 
element  similar  to,  but  much  more  active  than  chlorine.  In  this 
case  we  shall  see  that  fluorine  and  its  compounds  will  not  enter  into 
certain  reactions  analogous  to  those  shown  by  the  other  halogens, 
and  new  types  of  reactions  will  appear.  In  any  chemical  family 
the  element  having  the  lowest  atomic  weight  shows,  in  general, 
under  a  given  set  of  conditions,  the  greatest  activity.  The  differ- 
ence between  its  behavior  and  that  of  the  next  member  of  the  group 
is  much  greater  than  that  between  succeeding  members.  For 
this  reason  the  members  of  the  first  series  of  elements  in  the  peri- 
odic classification  are  not  as  typical  of  the  families  in  which  they 
occur  as  the  members  of  the  second  series  are.  In  general,  the 
groups  of  elements  which  show  the  most  striking  relationships  are 
composed  of  members  from  the  second  to  the  tenth  series. 

420.  Occurrence  and  Discovery. — Fluorine  occurs  in  small 
quantities  rather  widely  distributed.  It  is  present  in  several 
minerals,  the  most  important  of  which  are  fluorspar  or  fluorite, 
CaF2,  cryolite,  AlF3,3NaF,  and  apatite,  CaF2,3Ca3(PO4)2. 
Fluorine  also  occurs  in  traces  in  the  human  body,  and  is  present 
in  appreciable  amounts  in  the  enamel  of  the  teeth.  It  has  been 
suggested  that  when  the  teeth  of  growing  children  do  not  develop 
normally  one  cause  may  be  a  lack  in  the  food  of  sufficient  com- 
pounds containing  fluorine.  The  body  contains  a  large  number  of 
elements  in  small  quantities  and  we  must  obtain  these  from  food. 
Under  most  circumstances  an  adequate  supply  of  these  elements  is 


THE  HALOGEN  FAMILY  377 

obtained,  but  when  the  diet  is  restricted  to  a  few  substances  only 
there  may  result  a  deficiency  in  one  of  them. 

The  alchemists  were  familiar  with  the  fact  that  when  fluorspar 
is  heated  with  concentrated  sulphuric  acid  a  colorless  gas  is 
formed,  which  has  the  unusual  property  of  etching  and  dissolving 
glass.  Gay-Lussac  studied  the  reaction  in  1807  and  isolated  the 
gas  which  was  called  hydrofluoric  acid,  the  name  indicating  its 
preparation  from  fluorspar.  The  mineral  had  been  previously 
given  this  name  because  it  can  be  easily  melted — the  first  two  syl- 
lables in  fluorspar  being  derived  from  the  Latin  word  to  flow,  and 
the  last  being  the  general  name  for  minerals. 

For  over  seventy-five  years  chemists  attempted  in  vain  to  isolate 
from  fluorides  the  undiscovered  element  which  was  thought  to  be 
present.  All  attempts  were  unsuccessful  on  account  of  the  great 
activity  of  the  element,  which  was  called  fluorine.  Finally, 
Moissan  in  1886  came  to  the  conclusion  that  the  electric  current 
must  be  used  to  effect  the  decomposition  of  fluorides,  and  that  this 
must  be  carried  out  in  the  absence  of  water.  When  potassium 
fluoride  was  dissolved  in  liquid  hydrogen  fluoride  and  subjected 
to  the  action  of  an  electric  current,  a  light  yellow  gas  was  formed, 
which  proved  to  be  the  element  sought. 

421.  Preparation. — No  methods  similar  to  those  employed 
to  prepare  chlorine  can  be  used  for  the  isolation  of  fluorine  from  its 
compounds.  Oxidizing  agents  do  not  convert  hydrofluoric  acid 
into  fluorine  and  water,  because  fluorine  reacts  quantitatively 
with  water  to  form  hydrofluoric  acid  and  oxygen.  For  this 
reason,  also,  electrolysis  of  an  aqueous  solution  of  the  acid  does  not 
yield  the  free  halogen. 

Fluorine  is  prepared  in  the  way  used  by  Moissan.  The  appa- 
ratus in  which  the  electrolysis  is  carried  out  is  made  of  copper. 
Although  fluorine  attacks  the  metal,  it  soon  becomes  coated  with  a 
layer  of  copper  fluoride,  which  does  not  dissolve  in  the  hydro- 
fluoric acid  used  as  the  solvent  in  the  electrolysis  and  serves,  there- 
fore, as  a  protective  coating  for  the  metal.  The  electrodes  used 
are  made  from  an  alloy  of  platinum  and  iridium  as  all  other  sub- 
stances that  have  been  tried  are  attacked  by  the  nascent  fluorine. 
The  reaction  is  carried  out  at  a  low  temperature  in  order  to  reduce 
the  activity  of  the  halogen- 


378  INORGANIC  CHEMISTRY  FOR  COLLEGES 

422.  Physical  Properties. — Fluorine,  F2,  atomic  weight,   19, 
is  a  gas  with  a  pale  yellow  color;    it  has  been  converted  into  a 
liquid  which  boils  at  -186°  and  freezes  at  -223°. 

423.  Chemical  Properties. — Fluorine  reacts  with  all  elements 
except   chlorine,   oxygen,   nitrogen,   and   the   inert  gases   of  the 
helium  family.     When  perfectly  dry  it  does  not  attack  silicon 
dioxide  or  glass,  but  in  the  presence  of  a  trace  of  moisture  some 
hydrofluoric  acid  is  formed,  and,  as  a  result,  the  silicon  dioxide  is 
dissolved  and  the  glass,  which  contains  this  oxide,  disintegrated. 
It  is  for  this  reason  that  glass  vessels  cannot  be  used  in  preparing 
the  element.     The  noble  metals  are  converted  into  fluorides  very 
slowly  by  the  gas.     Fluorine  reacts  with  water  instantaneously; 
when  the  gas  is  passed  into  water  the  chief  products  of  the  reaction 
are  oxygen  and  hydrofluoric  acid,  which  we  shall  see  later  has 
the  formula  H2F2: 

2F2  +  2H2O  =  2H2F2  +  O2 

If  water-vapor  is  brought  into  contact  with  fluorine,  ozone  is 
formed.  On  account  of  the  difficulty  attending  the  preparation 
of  fluorine  and  the  great  activity  of  the  gas,  it  is  not  used  for  any 
purpose.  The  chief  interest  in  fluorine  centers  in  its  remarkable 
chemical  activity  and  the  relation  of  its  properties  to  those  of  the 
other  members  of  the  halogen  family. 

HYDROFLUORIC  ACID 

424.  Preparation. — The  formula  of  hydrofluoric  acid,  H2F2, 
indicates  that  it  is  a  dibasic  acid.     Like  acids  of  this  type  it  forms 
acid  salts,  one  of  which  has  the  formula  KHF2.     Many  such  acid 
salts  when  heated  are  broken  down  into  the  neutral  salt  and  the 
free  acid.     Advantage  is  taken  of  this  behavior  of  the  acid  fluoride 
to  prepare  the  anhydrous  acid: 

2KHF2  =  2KF  +  H2F2 

Hydrofluoric  acid  is  generally  used  in  solution,  and  is  com- 
monly prepared  by  heating  powdered  calcium  fluoride  with  con- 
centrated sulphuric  acid  in  a  retort  of  lead,  which  is  but  slightly 
attacked  by  the  gas;  the  latter  is  absorbed  in  water  contained  in  a 
bottle  made  of  hard  rubber  or  wax.  The  reaction  which  takes 
place  is  analogous  to  that  between  a  chloride  and  sulphuric  acid. 


THE  HALOGEN  FAMILY  379 

425.  Physical  Properties. — Hydrofluoric  acid  is  a  very  volatile, 
colorless  liquid  which  boils  at  19.4°.  It  forms  a  constant-boiling 
mixture  with  water,  which  boils  at  111°  at  750  mm.  pressure  and 
contains  43.2  per  cent  of  the  acid. 

The  molecular  weight  of  hydrogen  fluoride,  as  determined  by 
the  density  of  the  gas,  varies  markedly  with  the  temperature; 
the  volume  of  the  gas  at  26°  which  occupies  22.4  liters  when 
reduced  to  0°  and  760  mm.  weighs  51  grams;  a  similar  volume  of 
the  gas  at  a  temperature  above  90°  weighs  20  grams.  These  facts 
indicate  that  above  90°  the  molecular  weight  of  the  gas  is  20,  and 
its  formula  HF.  Below  90°  more  complex  molecules  are  present; 
the  molecular  weight  of  a  compound  having  the  formula  H2F2 
is  40  and  HaFa  is  60.  The  conclusion  is  drawn,  therefore,  that  in 
hydrogen  fluoride  these  three  types  of  molecules  are  present  in 
equilibrium:  H3F3  +±  H2F2  <=±  HF.  The  formula  H2F2  is  usually 
assigned  to  hydrofluoric  acid  because  the  molecular  weight  of  the 
acid  determined  by  the  freezing-point  method  corresponds  closely 
to  this  formula. 

When  two  or  more  molecules  of  the  same  kind  unite  to  form  a 
single  molecule  that  reverts  more  or  less  completely,  with  slight 
change  in  the  conditions,  to  the  molecules  from  which  it  was  formed, 
the  phenomenon  is  called  association.  The  vapor  of  hydrofluoric 
acid  from  a  temperature  just  above  the  boiling-point  of  the  liquid 
to  about  90°  is  associated  because  it  is  made  up  of  the  molecules 
HF,  H2F2,  and  HaFa  in  equilibrium.  It  will  be  recalled  that 
the  same  phenomenon  was  seen  in  the  case  of  nitrogen  tetroxide 
the  vapor  of  which  contains  two  kinds  of  molecules  in  equilibrium — 
N2C>4  <=±  2N02.  Associated  liquids  are  also  known;  it  is  highly 
probable  that  water  contains  different  kinds  of  molecules  in 
equilibrium — HeOs  +±  H4O2  +±  H2O.  Many  liquids  are  known 
the  properties  of  which  can  best  be  explained  on  the  assumption 
that  they  are  associated.  The  formulas  of  these  liquids  are  written 
in  such  a  way  as  to  indicate  their  relation  to  the  formulas  of  the 
simple  substances  of  which  they  are  made  up.  For  example,  in  the 
case  of  water  the  associated  molecules  are  usually  represented  by 
the  formulas  (H2O)s  and  (H2O)2  and  not  as  HeOs  and  H4O2  as 
given  above.  For  the  sake  of  simplicity  such  formulas  are  used 
only  when  the  fact  that  the  liquid  is  more  or  less  associated, 
is  an  important  factor  in  the  reaction  under  consideration. 


380  INORGANIC  CHEMISTRY  FOR  COLLEGES 

426.  Chemical  Properties  and  Uses. — When  hydrofluoric  ?cid 
is  dissolved  in  water,  the  acid  in  solution  is  much  weaker  than 
hydrochloric  acid.    At  18°  in  one-tenth  normal  solution,  the  latter 
is  dissociated  into  its  ions  to  the  extent  of  about  92  per  cent; 
under  the  same  conditions  hydrofluoric  acid  is  0.15  per  cent  dis- 
sociated.   This  fact  is  a  result,  in  all  probability,  of  the  great  activity 
of  fluorine,  but  the  exact  reason  is  not  apparent.     The  fact  that 
the  acid  exists  in  the  associated  form  is  also  traceable  to  the  same 
cause.     This  view  is  expressed  by  considering  the  complex  mole- 
cule as  made  up  as  indicated  in  the  following  graphic  formula: 
H  —  F  =  F— H.     From  this  point  of  view  we  would  not  expect  that 
the  ionization  of  this  complex  molecule  would  necessarily  bear  any 
relation  to  that  of  the  smaller  molecule  of  hydrochloric  acid, 
HC1. 

Hydrofluoric  acid  is  a  sufficiently  active  acid  to  react  with 
the  more  active  metals.  It  forms  fluorides  with  oxides  and 
hydroxides  of  the  metals.  It  differs  from  hydrochloric  acid  in 
being  dibasic  and,  accordingly,  forms  acid  salts  of  which  KHF2 
is  an  example. 

A  very  characteristic  property  of  hydrofluoric  acid  is  shown 
in  its  action  on  silicon  dioxide,  which  occurs  in  an  impure  form  as 
sand.  The  reaction  is  represented  by  the  following  equation: 

SiO2  +  2H2F2  =  SiF4  +  2H2O 

Silicon  fluoride  is  a  gas,  and,  as  a  result,  when  pure  silicon  dioxide 
is  treated  with  hydrogen  fluoride  in  the  form  of  a  gas  or  in  solution, 
the  oxide  is  converted  into  substances  which  are  volatile.  This 
reaction  is  utilized  in  quantitive  chemical  analysis  in  determining 
the  amount  of  silicon  present  in  substances  containing  the  element. 
The  latter  is  separated  as  silicon  dioxide  in  a  more  or  less  pure 
condition,  and  is  heated  until  its  weight  remains  constant;  the 
material  is  then  treated  with  a  solution  of  hydrofluoric  acid,  and 
heated  again  to  constant  weight;  the  loss  in  weight  equals  the 
weight  of  the  pure  silicon  dioxide  which  was  present,  and  from  it 
can  be  calculated  the  amount  of  silicon  in  the  compound  analyzed. 

427.  Salts  of  the  acids  containing  oxygen  may  be  considered 
as  made  up  of  oxides  of  metals  and  of  the  anhydrides  of  the  acids. 
Considered  in  this  way  the  formula  CaSiOs  may  be  written  as 


THE  HALOGEN  FAMILY  381 

follows:    CaO,SiO2.     When  the  salt  is  treated  with  hydrofluoric 
acid  both  oxides  are  converted  into  fluorides: 

CaO,SiO2  +  3H2F2  =  CaF2  +  SiF4  +  3H2O 


Common  glass  is  a  mixture  of  sodium  silicate,  NaoSiOa,  calcium 
silicate  and  silicon  dioxide.  When  it  is  treated  with  hydrofluoric 
acid,  the  silicon  passes  off  as  a  gas  in  the  form  of  the  tetrafluoride, 
SiF-i,  and  the  calcium  fluoride  and  sodium  fluoride  are  left  behind 
in  the  form  of  a  powder.  As  the  result  of  the  reaction  the  glass 
is  eaten  away,  or  etched,  wherever  hydrofluoric  acid  comes  in  con- 
tact with  it.  The  reaction  is  used  to  etch  designs  upon  glass. 
To  do  this  the  glass  is  covered  with  a  thin  layer  of  wax  and  the 
design  is  drawn  on  the  latter  with  a  sharp  point  which  cuts  through 
the  wax  to  the  glass.  When  exposed  to  hydrogen  fluoride  the 
uncovered  part  of  the  glass  is  attacked  by  the  acid,  and  when  the 
wax  has  been  removed,  the  design  is  clearly  seen  as  the  result  of 
the  fact  that  the  etched  surface  is  rough  and  semi-transparent  like 
ground  glass.  It  is  in  this  way  that  the  markings  are  put  on 
chemical  thermometers,  burettes,  and  other  measuring  instru- 
ments made  of  glass. 

428.  Properties  of  Fluorides.  —  The  fluorides  differ  markedly 
from  the  chlorides,  bromides,  and  iodides  in  solubility  and  other 
physical  properties.  Silver  fluoride  is  a  very  soluble  salt;  its 
molar  solubility  is  about  two  and  one-half  times  that  of  common 
salt.  On  the  other  hand,  calcium  fluoride  is  very  difficultly  soluble, 
whereas  the  other  halides  of  the  metal  dissolve  readily  in  water. 

The  fluorides  of  the  metals  are  very  stable  towards  heat; 
those  of  the  acid-forming  elements  react  with  water  in  a  way 
similar  to  that  in  which  the  analogous  compounds  of  the  other 
halogens  react. 

The  test  for  fluorides  is  based  on  the  action  of  hydrofluoric 
acid  on  glass.  The  substance  to  be  tested  is  ground  to  a  fine 
powder  and  mixed  with  concentrated  sulphuric  acid  in  a  lead  dish. 
A  glass  plate  coated  with  a  thin  layer  of  wax  or  paraffin,  in  which  a 
mark  has  been  scratched,  is  placed  over  the  dish.  After  a  few 
minutes  the  wax  is  removed  from  the  glass,  and  if  a  fluoride  was 
present  the  mark  will  appear  etched  into  the  glass. 


382  INORGANIC  CHEMISTRY  FOR  COLLEGES 

COMPOUNDS  OF  THE  HALOGENS  WITH  OXYGEN  AND  WITH 
HYDROGEN  AND  OXYGEN 

429.  Fluorine   forms   no    compound   with   oxygen,    although 
chlorine  and  iodine  do  form  such  compounds.     Acids  which  con- 
tain chlorine,  bromine,  or  iodine  in  combination  with    hydrogen 
and  oxygen  are  known;  they  form  well-characterized  salts,  some  of 
which    have   important   uses.     Attempts   to   prepare   analogous 
compounds  containing  fluorine  have  been  unsuccessful. 

430.  Chlorine  Monoxide. — When  chlorine  is  dissolved  in  water 
a  reaction  takes  place  as  the  result  of  which  hypochlorous  acid  and 
hydrochloric  acid  are  formed  (125) ; 

C12  +  H2O  =  HOC1  +  HC1 

If  sodium  hydroxide  is  added  to  the  solution,  or  if  chlorine  is 
passed  into  a  solution  of  the  base,  the  salts  of  the  two  acids  are 
formed.  Sodium  hypochlorite  is  an  unstable  substance.  When 
it  is  treated  with  sulphuric  acid,  the  hypochlorous  acid  set  free  de- 
composes into  its  anhydride  and  water: 

2NaOCl  +~H2SO4  =  Na2SO4  +  2HOC1 
2HOC1<=>H20  +  C120 

Chlorine  monoxide,  C120,  is  a  brownish-yellow  gas  (boiling- 
point  5°)  which  reacts  with  water  as  indicated  above.  It  is  best 
prepared  by  the  action  of  chlorine  on  freshly  precipitated  mercuric 
oxide,  which  has  been  carefully  dried: 

HgO  +  2C12  =  HgCl2  +  C120 

The  oxide  is  a  very  unstable  substance,  and  explodes  readily. 

431.  Hypochlorous  Acid. — We  have  just  learned  that  this  acid 
is  an  unstable  substance  and  decomposes  into  its  anhydride  and 
water  when  liberated  from  its  salts.     It  can,  however,  exist  in 
dilute  aqueous  solution.     It  decomposes  when  exposed  to  sun- 
light    (125),     2HOC1  =  2HC1  +  02,    and   when   brought   into 
contact  with  substances  that  can  be  oxidized,  and  for  this  reason 
it  is  used  as  a  bleaching  agent. 

Hypochlorous  acid  is  very  slightly  dissociated,  a  fact  which 
is  utilized  in  preparing  dilute  solutions  of  the  acid.  This  is  accom- 


THE  HALOGEN  FAMILY  383 

plished  by  passing  chlorine  into  a  dilute  solution  of  the  salt  of  a 
weak  acid  which  is  more  highly  dissociated  than  hypochlorous  acid. 
Sodium  carbonate  can  be  used,  for  example.  The  reactions 
involved  are  as  follows: 

2C12  +  H2O  +  [H2O]  <=±  2HOC1  +  [2HC1] 

Na2CO3  +  [2HC1]  =  2NaCl  +  [H2O]  +  C02 

Na2CO3  +  2C12  +  H2O  =  2HOC1  +  2NaCl  +  C02 

The  sodium  carbonate  reacts  with  the  hydrochloric  acid  formed  to 
produce  sodium  chloride,  but  does  not  appreciably  affect  the  weak 
hypochlorous  acid.  The  solution  so  produced  can  be  used  directly, 
if  the  presence  of  sodium  chloride  does  not  interfere  with  its  use, 
or  it  can  be  distilled;  if  it  is  dilute  the  distillate  obtained  is  a  dilute 
solution  of  hypochlorous  acid. 

Hypochlorous  acid  can  be  prepared  by  treating  bleaching 
powder  with  a  weak  acid.  When  bleaching  powder  (639)  is  dis- 
solved in  water  there  are  two  salts  present,  calcium  hypochlorite, 
Ca(OCl)2,  and  calcium  chloride,  CaCl2.  If  to  such  a  solution  an 
acid  is  added  stronger  than  hypochlorous  acid  and  weaker  than 
hydrochloric  acid,  the  former  will  be  set  free  and  the  latter  will 
not;  carbonic  acid  (CO2  +  H2O),  can  be  used  for  this  purpose: 

Ca++  +  2(001)"  +  2H+  +  C03"~  =  CaCO3  +  2HOC1 

Hypochlorous  acid  is  an  active  oxidizing  agent.  A  number  of 
reactions  have  already  been  considered  in  which  the  acid  acts  in 
this  way  (126).  It  oxidizes  iodine  in  aqueous  solution  to  iodic 
acid: 

5HOC1  +  I2  +  H2O  <=>  2HIO3  +  5HC1 

A  similar  reaction  takes  place  with  bromine,  and  bromic  acid, 
HBrO3,  is  formed.  It  is  for  this  reason  that  chlorine-water  is 
added  cautiously  in  testing  for  bromides  and  iodides;  an  excess  of 
the  former  would  convert  the  bromine  or  iodine  liberated  into 
bromic  acid  or  iodic  acid,  both  of  which  are  colorless,  and,  as  a 
result,  the  presence  of  the  halogens  would  be  overlooked. 

432.  Hypochlorites. — Most  of  the  salts  of  hypochlorous  acid 
are  unstable,  and  when  they  are  to  be  employed  for  any  purpose 


384  INORGANIC  CHEMISTRY  FOR  COLLEGES 

they  are  made  in  solution  and  used  in  this  form.     They  react  with 
hydrochloric  acid  and  form  chlorine: 

NaOCl  +  2HC1  =  NaCl  +  C12  +  H2O 

The  hypochlorite  of  sodium  is  prepared  by  passing  chlorine  into  a 
cold,  dilute  solution  of  sodium  hydroxide: 

C12  +  [H2O]  <=*  [HC1]  +  [HOC1] 
NaOH  +  [HC1]  =  NaCl  +  [H2O] 
NaOH  +  [HOC1]  =  NaOCl  +  H2O 

2NaOH  +  C12  =  NaCl  +  NaOCl  +  H2O 

433.  Chloric  Acid  and  Chlorates. — When  an  aqueous  solution 
of  sodium  hypochlorite  is  heated,  sodium  chlorate  is  formed. 
The  hypochlorites  of  other  metals  behave  in  a  similar  way.  The 
chlorate  commonly  prepared  and  used  is  the  potassium  salt;  on 
account  of  the  fact  that  its  solubility  is  relatively  small  in  cold 
water,  it  can  be  separated  from  the  potassium  chloride  formed 
along  with  it.  The  salt  can  be  made  by  passing  chlorine  into  a 
strong  hot  solution  of  potassium  hydroxide.  The  reactions  which 
take  place  can  be  analyzed  as  follows : 

3C12  +  [3H20]  =  [3HC1]  +  [3HC10] 
3KOH  +  [3HC1]  =  3KC1  +  3H2O 
3KOH  +  [3HC10]  =  [3KC1O]  +  [3H2O] 

[2KC1O]  =  2KC1  +  [20] 
[KC10]  +  [20]  =  KC103 


3C12  +  6KOH  =  KC1O3  +  5KC1+  3H2O 

After  completion  of  the  reaction  the  potassium  chlorate  is  obtained 
in  the  form  of  crystals  when  the  solution  is  cooled. 

This  method  of  preparing  potassium  chlorate  is  an  expensive 
one  on  account  of  the  fact  that  only  one-sixth  of  the  chlorine  and 
one-sixth  of  the  potassium  hydroxide  are  converted  into  the 
desired  salt;  the  rest  is  obtained  in  the  form  of  potassium  chloride, 
a  salt  that  occurs  in  nature  and  is,  accordingly,  less  expensive  than 
potassium  hydroxide. 

Potassium  chlorate  is  made  electrolytically  on  the  large  scale 
from  potassium  chloride.  When  a  current  is  passed  through  a 


THE  HALOGEN  FAMILY  385 

solution  of  the  chloride,  oxidation  takes  place  at  the  anode  and 
potassium  chlorate  is  formed;  hydrogen  is  set  free  at  the  cathode. 

434.  Chloric  acid,  HC1O3,  is  known  only  in  solution;  when  an 
attempt  is  made  to  concentrate  the  latter  it  decomposes  when 
about  40  per  cent  of  the  acid  is  present.     It  also  decomposes  when 
heated  above  40°.     For  these  reasons  chloric  acid  cannot  be  ob- 
tained from  its  salts  by  the  use  of  the  method  employed  in  the  case 
of  acids  which  boil  without  decomposition.     It  is  prepared  by 
double  decomposition  in  aqueous  solution  from  a  chlorate  and 
an  acid,  the  two  substances  being  selected  so  that  the  metal  in  the 
salt  and  the  radical  in  the  acid  are  those  of  an  insoluble  salt;  the 
latter  precipitates  and  the  desired  acid  remains  in  solution.     For 
example,  barium  chlorate  and  sulphuric  acid  could  be  used: 

Ba(ClO3)2  +  H2SO4  -  BaSO4  +  2HC1O3 

If  potassium  chlorate  is  treated  with  concentrated  sulphuric  acid, 
the  chloric  acid  set  free  decomposes  spontaneously  with  explosive 
violence;  among  the  products  formed  are  chlorine  dioxide,  C1O2, 
and  oxygen. 

Chloric  acid  is  an  active  oxidizing  agent;  a  strong  aqueous 
solution  of  the  acid  will  ignite  paper. 

435.  When  potassium  chlorate  is  heated  to  a  high  temperature 
it  loses  its  oxygen: 

2KC103  =  2KC1  +  302 

At  lower  temperatures  a  reaction  takes  place  which  resembles 
somewhat  the  change  that  occurs  when  hypochlorites  are  heated. 
In  the  case  of  the  chlorate  reduction  takes  place  simultaneously 
with  oxidation  and  a  perchlorate  is  formed: 

4KC1O3  =  KC1  +  3KC1O4 

Chlorates  are  active  oxidizing  agents.  Mixtures  made  up  of 
potassium  chlorate  and  readily  oxidizable  substances,  such  as  red 
phosphorus  and  sulphur,  explode  violently  when  struck.  Such 
mixtures  diluted  with  sand  or  fine  pebbles,  are  used  in  making  toy 
torpedoes  for  use  in  celebrating  Independence  Day. 

436.  Perchloric  Acid  and  Perchlorates. — Potassium  perchlorate 
is  about  one-twentieth  as  soluble  as  potassium  chloride  and  can 
be  readily  separated  by  crystallization  from  the  chloride  formed 


386  INORGANIC  CHEMISTRY  FOR  COLLEGES 

when  potassium  chlorate  is  cautiously  heated.  The  equation  for 
the  reaction  by  which  potassium  perchlorate  is  formed  has  been 
given  (435) .  Perchloric  acid  is  a  colorless  liquid  which  can  be  heated 
to  about  90°  without  change.  On  account  of  this  fact  and  the 
fact  that  it  boils  below  this  temperature  when  the  pressure  is 
reduced,  the  acid  can  be  prepared  by  treating  potassium  per- 
chlorate with  concentrated  sulphuric  acid  and  distilling  the 
acid  under  reduced  pressure.  The  process  has  to  be  carried  out 
with  great  care  and  the  pressure  must  be  as  low  as  possible  in 
order  that  distillation  may  take  place  at  a  low  temperature.  Per- 
chloric acid  has  been  recently  prepared  commercially  by  the 
electrolytic  oxidation  of  hydrochloric  acid. 

In  order  to  prevent  decomposition,  perchloric  acid  is  usually 
kept  in  solution  (70  per  cent  or  less).  The  acid  is  an  oxidizing 
agent,  but  not  so  active  as  the  other  oxygen  acids  of  chlorine. 
When  it  decomposes,  chlorine  dioxide  is  among  the  products 
formed. 

When  perchloric  acid  is  treated  at  a  low  temperature  with 
phosphorus  pentoxide,  which  is  a  powerful  dehydrating  agent,  the 
anhydride  of  the  acid  is  formed,  and  the  water  removed  from  tho 
the  acid  unites  with  the  pentoxide  to  form  metaphosphoric  acid : 

2HC104  +  P2O5  =  C1207  +  2HP03 

Perchloric  anhydride  is  obtained  from  the  mixture  by  careful 
distillation;  it  is  a  colorless  liquid,  which  boils  at  82°,  and  explodes 
when  struck  or  heated  above  its  boiling-point. 

The  salts  of  perchloric  acid  are  more  stable  than  the  chlorates, 
but  they  lose  their  oxygen  when  heated  and  give  it  up  readily  to 
substances  that  can  be  oxidized;  they  are  used  in  the  preparation 
of  various  kinds  of  fireworks. 

On  account  of  the  fact  that  potassium  perchlorate  is  a  difficultly 
soluble  salt,  perchloric  acid  is  used  in  the  quantitative  determina- 
tion of  potassium. 

437.  Chlorine  Dioxide.— When  chloric  acid  and  perchloric 
acid  decompose,  chlorine  dioxide,  C1O2,  a  yellow  gas  which  con- 
denses to  a  liquid  at  10°,  is  among  the  products  formed.  It 
explodes  violently,  and  a  large  amount  of  heat  is  evolved;  it  is  a 
highly  endothermic  compound. 


THE  HALOGEN  FAMILY  387 

Chlorine  dioxide  and  water  react  and  form  chlorous  and 
chloric  acids: 

2C1O2  +  H2O  =  HClOo  +  HC1O3 

The  chlorous  acid  formed  in  this  way  exists  only  in  solution. 
From  the  acid,  salts  called  chlorites  have  been  prepared. 

438.  Hypobromous  Acid  and  Bromic  Acid.  —  These  acids  and 
their  salts  are  well  known;    the  formulas  and  properties  of  the 
substances  are  similar  in  each  case  to  those  of  the  analogous 
chlorine  compounds.     The  acids  are  active  oxidizing  agents,  but 
less  so  than  hypochlorous  and  chloric  acid.     No  oxides  of  bromine 
have  been  isolated. 

439.  Oxygen  Acids  of  Iodine.  —  Hypoiodous  acid,  HIO,  and 
its  salts  are  known  in  solution,  but  they  pass  readily  into  iodic 
acid  and  iodates.     They  are  active  oxidizing  agents,  which  resem- 
ble bromic  acid  and  bromates  in  properties  and  modes  of  prepa- 
ration.    Sodium  iodate,   NalOa,   is  an  important   salt,   since   it 
occurs  in  Chile  saltpeter  and  is  the  chief  source  of  the  iodine  of 
commerce. 

Iodic  acid  can  be  readily  prepared  by  oxidizing  iodine  with 
strong  nitric  acid.  It  is  a  white  solid  which  loses  water  at  about 
170°  and  passes  into  iodic  anhydride: 

2HIO3  «=±  H20 


At  about  300°  the  anhydride  begins  to  decompose  into  iodine  and 
oxygen,  but  it  does  not  explode. 

440.  Compounds  of  the  Halogens  with  Themselves  and  with 
Other  Elements.  —  Two  chlorides  of  iodine,  the  monochloride,  IC1, 
and  the  trichloride,  Ida,  are  formed  as  the  result  of  the  direct 
union  of  the  halogens;  the  monochloride  is  a  red  crystalline  sub- 
stance and  the  trichloride  a  yellow  powder.  They  are  decom- 
posed by  water  and  by  heat.  Compounds  of  the  formula  IBr  and 
IFs  are  said  to  exist. 

The  halogens  form  compounds  with  other  acid-forming  ele- 
ments. Nitrogen  chloride,  NCls,  is  an  oily  liquid  which  is  formed 
when  a  solution  of  ammonium  chloride  is  treated  with  chlorine; 
it  is  highly  explosive.  No  bromide  of  nitrogen  is  known.  Nitro- 
gen iodide,  to  which  the  formula  Nl3,NHs  is  assigned,  is  prepared  by 
treating  a  strong  aqueous  solution  of  ammonia  with  iodine.  It 


388  INORGANIC  CHEMISTRY  FOR  COLLEGES 

separates  as  a  black  insoluble  precipitate,  which  decomposes  vio- 
lently when  touched  in  the  dry  condition. 

441.  The  Halogen  Family  and  the  Periodic  Law. — The  study 
of  the  properties  of  the  halogens  and  their  important  compounds 
has  brought  out  clearly  the  interesting  relationships  which  exist 
among  them.  In  general,  the  change  in  any  property  is  progressive 
when  we  pass  from  one  element  to  another  in  the  order  of  increasing 
atomic  weights.  If  the  property  is  a  physical  one  closely  associated 
with  weight,  such  as  density,  its  value  increases  with  increasing 
atomic  weight ;  if  it  is  not  so  associated,  the  change  may  be  in  the 
reverse  direction.  For  example,  the  solubilities  of  the  halides  of 
silver  decrease  as  we  pass  from  the  fluoride  to  the  iodide,  but  the 
solubilities  of  the  halides  of  sodium  increase  as  we  pass  in  the  same 
direction.  We  do  not  know  the  causes  underlying  solubility,  but 
the  fact  is  important  that  in  most  cases  there  is  a  progressive  change 
in  solubility  as  we  pass  from  one  compound  to  the  next  in  a  series 
of  salts  of  analogous  composition  derived  from  the  elements  in  a 
chemical  family. 

The  change  in  chemical  properties  is  also  progressive.  For 
example,  the  activity  of  the  halogens,  as  measured  by  the  energy 
liberated  when  they  unite  with  hydrogen,  decreases  with  increasing 
atomic  weight,  and  the  stability  of  the  compounds  formed  decreases 
in  the  same  order.  When  we  consider,  however,  the  compounds 
containing  oxygen,  and  oxygen  and  hydrogen,  the  order  is  reversed  ; 
the  most  stable  compounds  are  derived  from  iodine.  As  a  result 
of  these  facts  chlorine  will  drive  out  iodine  from  hydriodic  acid, 
whereas  iodine  will  convert  chloric  acid  into  iodic  acid.  The 
order  of  replacement  is  reversed  in  the  second  case  on  account  of 
the  fact  that  oxygen  compounds  containing  iodine  are  more 
stable  than  the  analogous  chlorine  compounds. 

According  to  the  periodic  table  the  elements  in  the  seventh 
group  have  the  valence  1  toward  hydrogen  and  7  toward  oxygen. 
The  valencies  indicated  in  the  table  are  the  highest  possible;  in 
many  cases  compounds,  the  possibility  of  the  existence  of  which  is 
indicated  by  the  law,  have  never  been  prepared.  Perchloric 
anhydride,  C^Oy,  is  the  only  oxide  in  which  a  halogen  atom  shows 
the  valence  7.  Valencies  between  1  and  7  are  shown  in  a  num- 
ber of  compounds.  The  graphic  formulas  of  some  of  these  are 
written  as  follows: 


THE  HALOGEN  FAMILY  389 

C1 


No     H— O— Cl=0     Cl/        H-O-C1/       H-O-C1^=O 
CK  ^O  ^0  \0 

Cl^O 


442.  The  Electronic  Theory  of  Valence. — A  very  brief  account 
has  already  been  given  of  the  recently  advanced  theory  of  the  con- 
stitution of  matter,  which  postulates  that  the  atoms  are  made  up  of 
charges  of  electricity  (398) .  According  to  this  view  the  unchange- 
able nucleus  of  the  atom  is  composed  of  positive  and  negative 
charges  of  electricity,  and  outside  of  these,  in  what  is  called  the 
outer  sphere,  there  are  a  definite  number  of  negative  charges  which 
come  into  play  when  the  atom  unites  with  another  atom.  These 
negative  charges  are  called  valence  electrons,  and  the  valence  of  an 
element  is  determined  by  the  number  of  these  electrons  possessed 
by  each  atom  of  the  element.  The  number  of  valence  electrons 
varies  from  0  in  the  case  of  the  inert  elements  in  group  0  of  the 
periodic  classification  to  8  in  the  case  of  the  elements  in  group  8, 
the  number  on  any  element  being  the  same  as  the  number  of  the 
group  in  which  it  occurs  in  the  classification. 

According  to  this  theory,  when  chemical  union  takes  place 
between  two  elements,  a  valence  electron  passes  from  one  element 
to  the  other,  and  there  is  established  between  the  two  a  so-called 
electrical  field  of  force,  which  holds  them  in  combination.  The 
element  which  loses  the  electron  is  the  more  positive  of  the  two. 
When,  for  example,  a  sodium  atom,  which  has  1  valence  electron 
(group  1),  unites  with  chlorine,  which  has  7  valence  electrons  (group 
7) ,  the  metallic  atom  loses  its  electron,  which  passes  to  the  chlorine 
atom.  This  is  expressed  in  the  electronic  formula  for  sodium 
chloride  as  follows:  Na  — >  Cl;  the  line  ordinarily  representing  a 
valence  bond  is  replaced  by  an  arrow  to  indicate  that  in  this  case 
an  electron  has  passed  from  the  sodium  atom  to  the  chlorine  atom. 
When  the  product  of  the  reaction,  sodium  chloride,  is  dissolved  in 
water,  the  latter  separates  the  atoms  from  each  other,  and,  as  a 


390  INORGANIC  CHEMISTRY  FOR  COLLEGES 

consequence,  the  sodium  which  in  its  union  with  chlorine  lost  a 
negative  charge,  becomes  positive,  a  sodium  ion,  and  has  the 
valence  +  1 ;  and  the  chlorine  which  has  gained  a  negative  charge 
becomes  negative  and  has  the  valence  —  1.  Losing  a  negative 
charge  of  electricity  is  the  same  as  gaining  a  positive  charge. 

On  account  of  the  fact  that  the  compounds  formed  as  the 
result  of  the  union  of  two  non-metallic  elements  do  not  dissociate 
into  ions,  the  union  between  such  elements  is  considered  to  be  of 
a  different  nature.  In  this  case  the  elements  are  said  to  share 
pairs  of  electrons  between  them.  In  chlorine  monoxide  the  union 
is  represented  thus :  Cl  :O  :C1.  The  two  dots  between  chlorine  and 
oxygen  indicate  that  each  element  furnishes  an  electron  which  is 
held  in  common  by  the  two  atoms.  There  is  not  a  sufficient  dif- 
ference in  chemical  properties  between  the  atoms  to  cause  the 
passage  of  an  electron  from  one  element  to  the  other  (806). 

It  will  be  seen  by  inspecting  the  periodic  table  that  there  is 
a  remarkable  fact  in  regard  to  the  positive  and  negative  valences  of 
all  the  elements.  In  group  7  the  elements  form  compounds  of  the 
type  HX,  in  which  the  valence  is  —  1  and  of  the  type  X2Oj  in  which 
the  valence  is  +  7.  In  group  6  the  valence  values  are  —  2  and  +  6, 
in  group  5  they  are  —  3  and  +  5,  and  in  group  4,'  —  4  and  +  4. 
The  sum  of  the  valencies  disregarding  the  signs  is  8  in  each  case. 
This  fact  is  a  help  in  remembering  the  valencies  of  elements;  if,  for 
example,  we  know  that  the  highest  valence  of  an  element  toward 
hydrogen  is  3,  then  we  know  that  its  valence  toward  oxygen  and 
other  negative  elements  is  5. 

The  facts  given  above  lead  to  the  conclusion  that  the  largest 
number  of  valence  electrons  which  can  exist  on  an  element  is  8. 
For  example,  chlorine,  which  has  7  valence  electrons,  can  receive 
but  1  more. 

EXERCISES 

1.  (a)  At  what  temperature  does  bromine  boil  when  the  pressure  on  it 
is  150  mm.?     (6)  What  weight  of  bromine  is  required  to  fill  at  20°  a  bottle 
which  holds  250  c.c.? 

2.  (a)  Write  an  equation  expressing  the  equilibrium  between  the  sub- 
stances present  in  a  solution  of  bromine  in  water.     (6)  Explain  what  happens 
if  air  is  bubbled  through  the  solution. 

3.  The  average  of  a  number  of  experiments  led  to  the  conclusion  that 
the  weight  of  1  liter  of  hydrogen  bromide  reduced  to  0°  and  760  mm.  is  3.612 
grams.     Using  1.008  as  the   atomic  weight  of  hydrogen  and    HBr    as  the 
formula  of  hydrogen  bromide  calculate  the  atomic  weight  of  bromine. 


THE  HALOGEN  FAMILY  391 

4.  (a)  State  what  would  happen  if  finely  divided  silver  is  shaken  with 
a  solution  of  hydrobromic  acid  and  bromine.     (6)  Devise  a  way  to  determine 
quantitatively  the  amount  of  each  in  the  solution. 

5.  What    (a)  weight  .and  what    (6)  volume  of  the  constant  boiling  aque- 
ous solution  of  hydrobromic  acid  are  required  to  react  with  10  grams  of  zinc? 

6.  The  reaction  between  sulphuric  acid  and  hydrobromic  acid  is  a  revers- 
ible one.     Under  what  conditions  could  it  be  used  to  prepare  a  solution  of 
hydrobromic  acid  from  bromine? 

7.  What  weight  of  iodine  must  be  dissolved  in  1000  c.c.  of  water  to  make 
(a)  a  0.1N  solution  and   (6)  a  0.1  molar  solution?     What  weight  of  Na2S2O3 
must  be  dissolved  in  100  c.c.  of  water  to  make    (c)  a  0.1N  solution  and   (d) 
a  0.1  molar  solution? 

8.  Forty  c.c.  of  a  0.05  molar  solution  of  I2  were  added  to  100  c.c.  of  an 
aqueous  solution  of  SO2.     The  excess  of  I2  required  15  c.c.  of  a  0.05  molar 
solution  of  Na2S2O3  to  react  with  it.     Calculate   (a)  the  weight  of  SO2  in  the 
solution,    (6)  percentage  by  weight  of  SO2  present  and    (c)  the  volume  of 
SO2  required  to  prepare  such  a  solution. 

9.  What  weight  of  iodine  is  required  to  make  100  c.c.  of  the  constant 
boiling  solution  of  hydriodic  acid? 

10.  Compare  the  action  of   (a)  light  and   (6)  heat  on  the  reactions  repre- 
sented by  the  following  equations : 

C12  +  H2O  +±  2HC1  +  O2,  and  I2  +  H2O  ±+  2HI  +  O2. 

11.  Devise  a  method  of  determining  quantitatively  the  amount  of  a  soluble 
fluoride  in  a  solution. 

12.  What  changes  in  color  are  observed  when  chlorine-water  is  added  to 
a  solution  containing  a  bromide  and  an  iodide?     (6)  Write  equations  for 
the  reactions  involved  including  those  which  lead  to  the  decolorization  of 
the  solution. 

13.  Devise  a  way  in  which  the  chlorine  set  free  could  be  determined  quanti- 
tatively when  a  solution  of  bleaching  powder  is  acidified. 

14.  How  could  you  distinguish  from  each  other  solutions  containing    (a) 
sodium  chloride,    (6)  sodium  hypochlorite,    (c)  sodium  chlorate,    (d)  sodium 
iodide,  and   (e)  sodium  bromide? 

15.  How  could  you  distinguish  the  following  from  one  another:   (a)  NaClO3, 
(6)  NaIO3,    (c)  NaF? 

16.  Compare  the  behavior  of  C1O2  and  NO2  when  brought  into  contact 
with  cold  water.    State  the  changes  in  valence  of  Cl  and  N  which  take  place 
in  the  two  cases. 


CHAPTER  XXVII 
SELENIUM  AND  TELLURIUM 

443.  Selenium  and  tellurium  are  in  the  sixth  group  in  the 
periodic  classification ,  in  which  oxygen  is  the  element  with  the 
lowest  atomic  weight;   they  form  with  sulphur  a  chemical  family 
in  which  the  relationships  between  the  members  of  the  family  are 
as  striking  as  in  the  case  of  the  halogens.     All  elements  in  the 
family  have  a  negative  valence  of  2  and  can  exhibit  a  positive 
valence  up  to  6.     Selenium  and  tellurium  form  compounds  anal- 
ogous   in    composition   to   £[28,  SO2,  H^SOs,   EfeSCU,  SC12,  etc. 
The  elements  are  present  in  the  earth  in  relatively  small  quantities 
only,  and  have   not   been  used  for  many  purposes.     They  and 
some  of  their  compounds  will  be  considered  briefly. 

444.  Occurrence. — Selenium  occurs  in  the  elementary  condi- 
tions associated  with  free  sulphur,  and  as  a  selenite  along  with  sul- 
phides.    It  occurs  in  some  varieties  of  pyrite,  FeS2.     It  was  dis- 
covered by  Berzelius  in  1817  in  the  flue  dust  separated  from  the 
gas  produced  by  burning  pyrite  in  the  manufacture  of  sulphuric 
acid. 

Tellurium  occurs  largely  in  combination  with  copper,  gold,  or 
silver  as  tellurides,  and  is  obtained  as  a  by-product  in  the  elec- 
trolytic refining  of  copper.  The  element  was  discovered  in  1783 
by  Miiller  von  Reichenstein  and  was  later  named  tellurium  on 
account  of  the  earthy  appearance  of  the  substance  from  which  it 
was  obtained  (tellus,  the  earth).  The  name  of  selenium,  discov- 
ered later,  was  derived  from  the  Greek  word  signifying  the  moon. 

445.  Preparation  and  Physical  Properties. — In  the  case  of  both 
selenium  and   tellurium  the  usual   method   of  preparation  is  the 
same,  namely,  the  precipitation  of  the  free  element  from  a  solution 
of  the  chloride  by  means  of  sulphur  dioxide: 

SeCU  +  2SO2  +  4H2O  =  Se  +  2H2SO4  +  4HC1 
392 


SELENIUM  AND  TELLURIUM  393 

The  elements  can  be  separated  by  utilizing  the  fact  that  tellurium 
is  not  precipitated  in  concentrated  hydrochloric  acid  solution, 
while  selenium  is  precipitated. 

Selenium  prepared  in  this  way  is  a  red  powder  and  tellurium  a 
black  powder.  The  elements  exist  in  allotropic  forms  (see  sul- 
phur, 273).  When  selenium  is  melted  and  allowed  to  solidify  it  is 
converted  into  an  amorphous  solid,  which  melts  at  217°,  resembles 
lead  in  color,  and  has  some  of  the  physical  properties  of  metals.  It 
conducts  electricity  to  a  slight  extent,  the  conductivity  being 
markedly  influenced  by  exposure  to  light.  This  fact  has  been 
ingeniously  utilized  in  the  invention  of  a  photometer,  which  has 
been  used  to  measure  the  relative  intensities  of  the  light  given  off  by 
different  stars.  Selenium  is  used  to  some  extent  in  coloring  glass 
and  enamels;  it  gives  to  them  a  characteristic  pink  or  red  color. 

Tellurium  melts  at  452°  and  on  solidification  is  obtained  as 
a  crystalline  substance,  which  resembles  silver  in  color.  Its 
metallic  properties  are  more  marked  than  are  those  of  selenium. 
It  has  no  uses  at  present. 

446.  Chemical  Properties. — Selenium  and  tellurium  burn  in  the 
air  and  form  dioxides,  SeO2  and  TeO2.     They  also  unite  with 
chlorine  and  form  compounds  represented  by  the  following  for- 
mulas:   Se2Cl2,  SeCU,  TeCl2,  TeCU.     Both  elements  react  when 
heated  with  metals,  and  selenides  and  tellurides  are  formed. 

447.  Hydrides. — When  ferrous  selenide  is  treated  with  con- 
centrated hydrochloric  acid,  hydrogen  selenide  is  formed: 

FeSe  +  2HC12  =  FeCl2  +  H2Se 

The  compound  is  a  poisonous  gas,  which  has  a  disagreeable, 
characteristic  odor;  it  dissolves  in  water  and  shows,  in  general,  a 
chemical  behavior  similar  to  that  of  hydrogen  sulphide.  Hydrogen 
telluride  is  made  in  a  similar  way  and  shows  similar  properties. 
Aqueous  solutions  of  the  hydrides  of  sulphur,  selenium,  and  tellu- 
rium are  decomposed  slowly  by  the  oxygen  in  the  air,  as  the  result 
of  which  water  and  the  free  elements  are  formed ;  the  rate  at  which 
this  decomposition  takes  place  increases  with  increasing  atomic 
weight  of  the  element.  As  we  pass  from  sulphur  to  tellurium 
metallic  properties  become  more  evident  and,  as  a  result,  the 
hydrides  become  less  stable.  The  selenides  and  tellurides  of  the 


394  INORGANIC  CHEMISTRY  FOR  COLLEGES 

metals,  with  the  exception  of  those  of  the  alkali  metals,  are  insol- 
uble in  water. 

448.  Oxides  and  Oxygen  Acids. — The  dioxides  of  selenium  and 
tellurium  are  formed  when  the  elements  are  burned  in  the  air  or 
oxidized  with  nitric  acid.  Selenium  dioxide  is  a  white  solid,  which 
sublimes  readily  and  forms  long  needles;  it  is  soluble  in  water  and 
reacts  with  it  to  some  extent  to  form  selenous  acid: 

SeO2  +  H20  <=±  H2SeO3 

The  salts  of  selenous  acid,  the  selenites,  resemble  closely  the  sul- 
phites. Selenous  acid  is  reduced  by  sulphurous  acid  to  selenium: 

H2SeO3  +  2H2SO3  =  Se  +  2H2SO4  +  H2O 

This  reaction  is  commonly  used  in  testing  for  selenium  because  all  of 
its  compounds  can  be  readily  converted  into  selenous  acid,  and  the 
formation  of  the  red  precipitate  is  characteristic.  It  can  also  be 
used  for  the  quantitative  determination  of  the  element  because 
selenium  can  be  heated  below  100°  without  change. 

Tellurium  dioxide  is  insoluble  in  water,  but  dissolves  in  solu- 
tions of  alkalies  and  forms  salts  of  tellurous  acid,  H2TeO3,  which  is 
precipitated  when  soluble  tellurites  are  treated  with  an  acid. 
Tellurium  dioxide  shows  weakly  basic  properties;  it  dissolves 
in  concentrated  hydrochloric  acid  and  forms  a  chloride,  TeCU, 
which,  however,  is  decomposed  when  water  is  added  to  the  solu- 
tion: 

TeCU  +  3H2O  <=±  H2Te03  +  4HC1 

It  will  be  recalled  that  the  chlorides  of  the  strongly  metallic  ele- 
ments are  not  decomposed  by  water,  and  that  those  of  the  acid- 
forming  elements  are  decomposed,  forming  hydrochloric  acid  and 
the  acid  derived  from  the  acid-forming  elements.  It  is  seen  from 
the  above  that  tellurium  chloride  stands  between  these  two 
extremes;  it  is  more  or  less  decomposed,  depending  upon  the  con- 
centration of  hydrochloric  acid  present. 

Tellurium  dioxide  dissolves  in  concentrated  sulphuric  acid, 
and  the  sulphate  formed  has  the  composition  represented  by  the 
formula  [TeO2]2,SO3.  The  normal  sulphate  of  tellurium  would 
have  the  formula  TeO2,2S03;  it  is  seen,  therefore,  that  the  com- 
pound formed  contains  a  large  excess  of  the  oxide  of  the  basic 


SELENIUM  AND  TELLURIUM  395 

element  over  that  present  in  the  normal  salt.  The  compound 
formed  is  a  basic  sulphate  (241).  In  general,  weakly  basic  ele- 
ments form  basic  salts.  The  oxide  forms  a  basic  nitrate  when 
dissolved  in  concentrated  nitric  acid;  its  formula  is  generally 
written  Te2Os(OH)NO3.  It  can  be  considered  as  made  up  as 
represented  by  the  following  formula:  [TeO2]4,H2O,N205. 

Tellurium  dioxide  and  tellurous  acid  react  with  bases  and  form 
salts  in  which  tellurium  plays  the  part  of  the  acid-forming  element; 
they  also  react  with  acids  and  form  salts  in  which  tellurium 
plays  the  part  of  a  base-forming  element.  Other  elements  form 
compounds  which  behave  in  this  way  and  to  emphasize  this  prop- 
erty such  compounds  are  said  to  be  amphoteric,  the  adjective 
being  derived  from  the  Greek  word  meaning  both. 

Selenic  acid,  H2SeO4,  can  be  prepared  by  oxidizing  selenous 
acid  with  chlorine-water: 

H2SeO3  +  C12  +  H2O  +±  H2SeO4  +  2HC1 

The  reaction  is  a  reversible  one,  for  in  the  presence  of  strong  hydro- 
chloric acid  selenic  acid  is  reduced  to  selenous  acid.  Selenic  acid 
is  a  more  active  oxidizing  agent  than  sulphuric  acid,  which,  it  will 
be  recalled,  will  not  oxidize  hydrochloric  acid;  it  will  oxidize  gold, 
which  is  not  attacked  by  nitric  acid. 

Telluric  acid  is  prepared  by  treating  tellurous  acid  with 
chromic  acid,  which  is  a  very  active  oxidizing  agent.  The  formula 
of  the  acid  is  HeTeOe,  and  is  written  in  this  way  and  not  as 
H2TeO4,2H2O  in  order  to  bring  out  the  fact  that  the  water  is  not 
combined  as  it  is  in  a  salt  containing  water  of  crystallization 
(water  of  hydration).  The  acid  is  a  white  crystalline  solid,  which 
when  heated  loses  water  and  is  converted  into  tellurium  trioxide, 
TeOa,  or,  at  higher  temperatures,  into  the  dioxide,  TeO2.  Telluric 
acid  is  readily  reduced  when  heated  with  a  solution  of  hydrochloric 
acid,  and  the  tellurous  acid  formed  is  converted  into  tellurium 
tetrachloride. 

A  careful  comparison  of  the  behavior  of  the  compounds  of 
selenium  and  tellurium  with  that  of  the  analogous  compounds  of 
sulphur,  will  bring  out  clearly  the  relationships  in  physical  and 
chemical  properties  which  occur  in  a  well-characterized  family  of 
elements. 


CHAPTER  XXVIII 
PHOSPHORUS,  ARSENIC,  ANTIMONY,  AND  BISMUTH 

449.  The  elements  to  be  considered  in  this  chapter  are  mem- 
bers of  the  fifth  group  in  the  periodic  classification.  They  show 
in  their  physical  properties  and  chemical  behavior  a  wide  varia- 
tion, which,  however,  is  systematic  and  progressive  with  changing 
atomic  weight. 

There  is  no  striking  similarity  between  nitrogen,  the  first  ele- 
ment in  the  group,  and  phosphorus,  either  in  the  elementary  con- 
dition or  in  their  compounds.  Nitrogen  is  inert  except  at  high 
temperatures;  when  it  unites  with  oxygen  the  reaction  is  endo- 
thermic,  and  the  compounds  formed  decompose  more  or  less 
readily  with  the  evolution  of  energy;  nitric  acid,  as  a  result,  is  an 
active  oxidizing  agent.  Phosphorus,  on  the  other  hand,  is  one  of 
the  most  active  elements  at  ordinary  temperatures;  it  takes  fire 
spontaneously  in  the  air,  and  when  it  burns  a  large  amount  of 
heat  is  given  off;  phosphoric  acid  and  its  anhydride  are  very  stable 
substances  and  do  not  act  as  oxidizing  agents.  It  has  already 
been  pointed  out  that  in  most  cases  the  first  member  in  a  group  in 
the  periodic  classification  differs  markedly  in  chemical  behavior 
from  the  other  members  of  the  group,  and  the  fact  was  illustrated 
in  the  case  of  fluorine.  The  difference  between  nitrogen  and 
phosphorus  is  much  greater  than  that  between  fluorine  and 
chlorine.  In  the  case  of  the  halogen  family  there  is  a  progressive 
change  in  activity  of  the  elements  as  measured  by  the  amount 
of  energy  set  free  when  they  unite  with  another  element.  A 
similar  change  in  activity  is  exhibited  by  phosphorus,  arsenic, 
antimony,  and  bismuth,  and  they,  accordingly,  constitute  a  chem- 
ical family. 

The  combining  power,  however,  of  nitrogen  and  phosphorus 
as  measured  by  valence,  is  the  same,  since  both  elements  have  the 
valencies  normal  to  all  members  of  the  fifth  group,  namely  +5 

396 


PHOSPHORUS,  ARSENIC,  ANTIMONY   AND  BISMUTH     397 

and  —3;  as  a  result,  the  compounds  of  the  two  elements  have 
similar  formulas  such  as  NH3,  PH3,  N2O3,  P2O3,  HNO3,  HPO3, 
etc.  But  a  similarity  in  the  chemical  formulas  of  two  compounds 
is  not  necessarily  associated  with  similarity  in  chemical  behavior. 
The  energy  relationships  are  the  determining  factor  in  the  latter, 
and,  as  has  been  said,  nitrogen  and  phosphorus  differ  markedly  in 
this  respect.  For  these  reasons  phosphorus  is  considered  the  first 
member  of  the  chemical  family  in  the  fifth  group  in  the  periodic 
classification  of  the  elements. 

PHOSPHOKUS 

450.  Occurrence. — Phosphorus  occurs  widely  distributed  over 
the  earth's  surface  in  the  form  of  phosphates  in  the  soil,  and  is 
present  in  the  so-called  phospho-proteins,  which  are  present  in  all 
living  things.     The  bones  and  teeth  of  animals  contain  approxi- 
mately 30  per  cent  of  calcium  phosphate.     The  element  occurs  in 
several   minerals,   of    which    calcium    phosphate    (phosphorite), 
Ca3(PO4)2,  and  apatite,  CaF2,3Ca3(P04)2,  are  examples.     Large 
beds    of    phosphate    rock   are    found    in    the    United  States  in 
Florida,  Georgia,    Tennessee,  North  and  South  Carolina,  Utah, 
Montana,  and  Wyoming;  they  also  occur  in  Ontario,  Tunis,  and 
Algeria. 

451.  Discovery  and  Preparation. — Phosphorus  was  discovered 
in  1669  by  Brand,  an  alchemist,  who  lived  in  Hamburg,  Ger- 
many.    In  his  search  for  the  philosopher's  stone  he  distilled  at  a 
high  temperature  the  residue  obtained  by  evaporating  urine  to 
dryness.     The   properties   of   the   substance   were   such   that   it 
aroused  a  great  deal  of  interest,  and  it  was  exhibited  in  the  courts 
of  Europe.    When  phosphorus  is  spread  over  any  surface  it  glows  in 
the  dark  with  a  dull  light,  which  constantly  shifts  in  intensity  from 
place  to  place.     It  was  this  property  of  the  element  that  led  to  its 
name,  which  was  derived  from  the  Greek  words  signifying  light- 
bearer. 

The  method  of  preparing  phosphorus  from  calcium  phosphate 
obtained  from  bones  was  first  published  by  Scheele  in  1771.  In 
this  method,  which  was  used  for  many  years  to  prepare  the  element 
on  the  commercial  scale,  calcium  phosphate  is  heated  with  sul- 
phuric acid  and  a  small  quantity  of  water;  the  calcium  sulphate 


398  INORGANIC  CHEMISTRY  FOR  COLLEGES 

formed  is  separated  by  filtration,  and  the  residue,  which  con- 
tains phosphoric  acid,  is  evaporated  and  distilled  at  a  very  high 
temperature  with  carbon  in  an  earthenware  retort.  The  phos- 
phorus, which  passes  over  as  a  gas,  is  collected  under  water  and  is 
obtained  as  a  yellow  solid. 

In  the  industrial  method  of  preparation  now  used  the  reac- 
tion is  carried  out  in  one  operation  in  an  electric  furnace,  as  a 
very  high  temperature  (white  heat)  is  required.  The  calcium 
phosphate  is  decomposed  by  silicon  dioxide  (sand)  which  replaces 
the  sulphuric  acid  used  in  the  older  method;  and  the  phosphorus 
pentoxide  liberated  as  a  result  is  reduced  to  phosphorus  by  the 
carbon  present.  The  reactions  involved  are  indicated  by  the  fol- 
lowing equations  in  which  calcium  phosphate,  Cas(PO4)2,  and 
calcium  silicate,  CaSiOs,  are  represented  as  made  up  in  each  case  of 
basic  and  acidic  oxides: 

(CaO)3,P2O5  +  3SiO2  =  3(CaO,SiO2)  +  P2O5 
P2O5  +  5C  =  2P  +  5CO 

The  ^apor  from  the  furnace  is  cooled  and  led  into  water  where  the 
phosphorus  condenses  to  a  solid.  To  obtain  it  in  a  convenient 
form  it  is  melted  under  water  and  cast  in  molds  into  sticks. 

452.  Physical  Properties. — Phosphorus  exists  in  three  allo- 
tropic  modifications,  which  differ  markedly  in  physical  properties. 
White  phosphorus  obtained  when  the  vapor  of  the  element  is 
condensed,  is  a  horny  substance  which  can  be  cut  with  a  knife; 
it  has  the  density  1.83,  melts  at  44°,  boils  at  287°,  and  is  very  sol- 
uble in  carbon  disulphide,  and  less  so  in  ether,  chloroform,  and 
other  organic  liquids.  Commercial  white  phosphorus  has  a  light 
yellow  color,  which  is  due  to  traces  of  impurities. 

Up  to  about  500°  the  vapor  of  phosphorus  has  a  density  which 
leads  to  the  conclusion  that  its  molecule  contains  four  atoms.  At 
1700°  the  molecules  are  partly  diatomic,  P2;  the  lowering  of  the 
freezing-point  of  solutions  of  phosphorus  indicates  that  in  this  con- 
dition the  molecule  is  tetra-atomic,  P4. 

Red  phosphorus  is  a  red  powder,  which  consists  of  microscopic 
crystals.  It  is  prepared  by  heating  the  white  variety  at  about 
250°.  It  is  insoluble  in  all  liquids,  and  does  not  melt  when  heated, 
but  changes  into  vapor  which  is  identical  with  that  obtained  from 


PHOSPHORUS,  ARSENIC,  ANTIMONY,  AND  BISMUTH     399 

white  phosphorus.  The  density  of  red  phosphorus  varies  between 
2.05  and  2.34,  and  is  determined  by  the  conditions  which  exist 
when  it  is  made. 

Black  phosphorus  has  been  obtained  by  heating  the  white 
variety  at  200°  under  a  pressure  of  1200  kilograms  per  square 
centimeter,  which  is  approximately  116  atmospheres.  It  is  a 
black,  lustrous  substance,  which  has  the  density  2.69,  and  con- 
ducts electricity. 

453.  Chemical  Properties. — White  phosphorus  reacts  so  readily 
with  the  oxygen  of  the  air  that  it  is  kept  under  water.  Its 
kindling-point  in  the  air  is  about  35°,  and  as  it  oxidizes  freely  at 
lower  temperatures,  it  is  apt  to  take  fire  spontaneously. 

Phosphorus  burns  in  chlorine  and  forms  the  pentachloride, 
PCls,  and  by  regulating  the  conditions  the  trichloride,  PCls, 
can  be  formed.  It  unites  vigorously  also  with  the  other  halogens, 
and  combines  with  most  elements  when  heated  with  them;  with 
metals,  phosphides  are  formed,  of  which  calcium  phosphide, 
CasP2,  is  an  example. 

If  white  phosphorus  ignites  while  in  contact  with  the  flesh  it 
produces  a  severe  burn,  which  heals  very  slowly;  it  should  not  be 
touched,  therefore,  and  should  be  handled  with  great  caution  with 
the  aid  of  pincers.  The  element  is  an  active  poison,  about  0.15 
gram  being  a  fatal  dose  for  man.  It  is  used  on  account  of  this 
property  as  an  ingredient  of  rat  poisons,  which  contain  usually,  in 
addition  to  the  element,  a  fat  and  a  diluent  such  as  flour.  White 
phosphorus  was  formerly  much  employed  in  making  matches,  but 
as  the  workmen  in  the  factories  often  developed  a  terrible  disease 
which  attacked  the  teeth  and  jaw-bone,  a  prohibitive  tax  was  put 
on  the  matches  manufactured  in  this  way.  Red  phosphorus, 
which  is  not  poisonous,  is  now  used  for  this  purpose. 

When  phosphorus  burns,  the  pentoxide  formed  is  produced  as  a 
dense  white  cloud,  the  obscuring  power  of  which  is  greater  than 
that  of  a  cloud  produced  in  any  other  way.  For  this  reason  shells 
containing  phosphorus  were  used  in  the  recent  war  in  setting 
up  a  cloud  barrage,  and  in  determining  the  accuracy  of  gun-fire. 
The  puff  of  white  smoke  produced  when  a  shell  containing  phos- 
phorus exploded  could  be  seen  at  a  long  distance,  and  the  place 
where  it  fell  could  be  accurately  noted.  White  phosphorus  was 
also  used  in  tracer  bullets  to  ignite  the  hydrogen  in  war-balloons. 


400  INORGANIC  CHEMISTRY  FOR  COLLEGES 

When  red  phosphorus  is  formed  by  heating  the  white  variety 
of  the  element,  the  change  is  accompanied  by  the  evolution  of  a 
large  amount  of  heat;  the  red  form  of  the  element  is  much  less 
active,  as  we  might  expect,  than  the  white.  White  phosphorus 
slowly  changes  in  the  light  to  the  red  form;  it  is  for  this  reason 
that  sticks  of  the  element  in  the  laboratory  become  colored  on  the 
outside,  the  shade  being  dependent  on  the  intensity  of  the  light 
and  the  time  the  element  has  been  exposed. 

The  kindling  temperature  of  red  phosphorus  in  air  is  about  240°. 
Red  phosphorus,  which  is  not  poisonous,  enters  into  reactions  with 
other  elements  to  form  the  same  compounds  produced  from  the 
white  variety.  Most  of  the  phosphorus  made  is  used  in  the  manu- 
ufacture  of  matches.  (See  problem  15,  page  420.) 

454.  Phosphorescence. — The  slow  oxidation  of  phosphorus  in 
the  air  is  accompanied  by  the  giving  off  of  a  feeble  light;  the 
element  is  said  to  phosphoresce.  The  chemical  change  which 
takes  place  involves,  in  all  probability,  the  formation  at  first  of 
phorphorus  trioxide,  and  the  subsequent  oxidation  of  the  latter 
to  the  pentoxide.  This  appears  to  be  a  reasonable  explanation 
in  view  of  the  fact  that  the  vapor  of  phosphorus  trioxide  phos- 
phoresces in  the  air.  The  phenomenon  occurs  during  the  slow 
oxidation  of  phosphorus  at  room  temperature,  only  when  the 
pressure  of  the  oxygen  in  contact  with  the  element  is  200  mm.  or 
less. 

465.  Other  substances  phosphoresce  and  certain  living  things  emit  a 
light  which  resembles  that  given  off  by  phosphorus;  among  the  latter  are 
glow-worms,  fire-flies,  and  certain  varieties  of  fish.  The  phosphorescence 
frequently  observed  in  the  wake  of  a  steamer  at  sea  is  produced  as  the  result 
of  the  agitation  of  small  microscopic  animals.  The  phenomenon  of  phos- 
phorescence is  an  important  one  about  which  little  is  known.  It  is  produced, 
no  doubt,  as  the  result  of  the  direct  change  of  chemical  energy  into  light. 
When  substances  burn  with  the  evolution  of  light,  the  latter  is  produced 
as  the  result  of  the  fact  that  the  heat  liberated  in  the  chemical  reactions  sets 
the  molecules  in  such  rapid  vibration  that  they  emit  light.  The  wave- 
length of  the  light  emitted  where  a  substance  is  heated  is  determined  by  the 
temperature  attained  and  in  most  cases  is  the  same  whatever  the  nature  of 
the  substance.  The  radient  energy,  a  part  of  which  is  light,  given  off  from 
a  glowing  body  is  proportional  approximately  to  the  fourth  power  of  the 
temperature  (Stephan's  law). 

There  are  certain  substances  the  behavior  of  which  is  not  in  accord  with 
this  law.  This  is  notably  the  case  with  phosphorus,  which  emits  light  at 
ordinary  temperatures,  and  with  magnesium.  When  the  latter  burns  an 


PHOSPHORUS,  ARSENIC,  ANTIMONY,  AND  BISMUTH     401 

intense  white  light  is  given  off  which  is  utilized  in  photography,  magnesium 
being  one  of  the  ingredients  of  flash-powders.  The  temperature  produced 
when  the  metal  burns  is  over  3000°  below  that  at  which  a  body  would  give 
off  light  of  this  character  if  it  were  the  result  of  heat  energy  alone.  In  such 
cases  as  these  a  part  of  the  chemical  energy  is  directly  changed  into  light 
energy.  This  kind  of  change  is  utilized  when  gas  is  burned  in  connection 
with  a  Welsbach  gas-mantle.  The  materials  of  which  the  mantle  is  made 
apparently  facilitate  the  change  of  chemical  energy  into  visible  light,  for 
the  temperature  of  the  mantle  is  far  below  that  necessary  to  produce  the 
quality  of  light  obtained.  The  extent  of  this  change  is  very  small,  however, 
compared  with  that  of  the  change  of  chemical  energy  into  heat. 

456.  Oxides  of  Phosphorus. — When  phosphorus  is  burned  in 
air  or  oxygen,  phosphorus  pentoxide,  P2Os,  is  formed.     The  com- 
pound is  a  white  powder,  which  can  be  melted.     It  reacts  with 
water  violently  and  is  converted  into  metaphosphoric  acid: 

P2O5  +  H2O  =  2HPO3 

On  account  of  this  property  phosphorus  pentoxide  is  used  as  a 
drying  agent  for  liquids  and  gases  with  which  it  does  not  react. 
It  is  the  most  efficient  agent  known  for  this  purpose.  The  oxide 
reacts  with  many  compounds  containing  hydrogen  and  oxygen, 
and  extracts  from  them  these  elements  in  the  proportion  required 
to  form  water;  it  is  thus  a  dehydrating  agent.  Its  behavior  in 
this  way  with  nitric  acid  and  sulphuric  acid  has  been  noted; 
in  each  case  the  anhydride  of  the  acid  is  formed,  along  with  meta- 
phosphoric acid.  The  pentoxide  will  remove  hydrogen  and  oxygen 
from  paper,  which  is  made  of  cellulose,  (CeHioOs)*,  and  the  carbon 
left  behind  will  be  evident  as  the  result  of  the  black  color  produced 
when  the  oxide  came  in  contact  with  the  paper. 

When  phosphorus  pentoxide  is  heated  with  water,  the  meta- 
phosphoric acid  first  formed  is  converted  into  phosphoric  acid, 
HsPC^;  for  this  reason  the  pentoxide  is  called  phosphoric  anhy- 
dride. 

457.  The  trioxide  of  phosphorus,  P2Os,  is  formed  when  the 
element  is  oxidized  in  a  limited  supply  of  air.     It  is  a  white  solid 
which  phosphoresces  in  air,  and  oxidizes  to  the  pentoxide.     It 
melts  at  22.5°  and  boils  at  173°.     Phosphorus  trioxide  reacts  with 
water  at  ordinary  temperatures  to  form  phosphorous  acid,  HsPOs. 

458.  The  Phosphoric  Acids. — Three  well-characterized  acids 
are  derived  from  the  single  anhydride,  P205.     Since  they  are 


402  INORGANIC  CHEMISTRY  FOR  COLLEGES 

closely  related  and  each  contains  phosphorus  with  the  valence  5, 
they  are  called  phosphoric  acids,  and  a  prefix  is-  added  to  the  name 
to  designate  each  one.  Metaphosphoric  acid  has  the  formula 
HPO3(P205,H2O),  pyrophosphoric  acid  H4P2O7(P205,2H2O),  and 
orthophosphoric  acid  H3PO4(P2O5,3H2O).  The  relationship  exist- 
ing between  the  acids  is  clearly  brought  out  by  indicating  their 
composition  in  the  way  shown  in  the  parentheses.  When  phos- 
phorus pentoxide  is  dissolved  in  water,  metaphosphoric  acid  is  first 
formed.  On  standing  or  on  heating  the  solution,  further  hydration 
takes  place  and  orthophosphoric  acid  is  produced.  When  the 
latter  is  obtained  by  evaporating  the  water  and  is  heated  for  some 
time  at  255°  it  is  converted  into  pyrophosphoric  acid.  Meta- 
phosphoric acid  is  called  glacial  phosphoric  acid  in  commerce,  and 
is  usually  obtained  in  the  form  of  transparent  sticks. 

459.  Orthophosphoric  Acid. — The   calcium  salt  of  this  acid 
occurs  in  nature  and  is  used  as  the  commercial  source  of  the  acid, 
which  is  prepared  by  treating  the  mineral  with  sulphuric  acid. 
The  acid  may  be  prepared  by  oxidizing  red  phosphorus  with  strong 
nitric  acid.     It  is  a  white  solid  which  melts  at  42.3°. 

Orthophosphoric  acid,  HsPO^  contains  3  hydrogen  atoms 
which  can  be  replaced  by  metallic  atoms;  it  is,  consequently, 
tribasic.  When  dissolved  in  water  it  ionizes  as  indicated  by  the 
following  formulas : 

H3P04  +±  H+  +  H2P04~  <=*  H+  +  HP04— '  ±5  H+  +  PO4~ 

In  0.1N  solution  the  first  step  in  the  ionization  takes  place  to  the 
extent  of  about  27  per  cent  at  18°;  the  bivalent  and  trivalent  ions 
are  formed  to  a  much  smaller  extent.  Orthophosphoric  acid  is, 
accordingly,  a  much  weaker  acid  than  sulphuric  acid,  which  is 
ionized  to  the  extent  of  61  per  cent  at  the  same  concentration. 

460.  Orthophosphates. — Orthophosphoric  acid,  like  other  tri- 
basic acids,  forms  three  classes  of  salts  which  are  produced  as  the 
result  of  the  replacement  of  one-third,  two-thirds,  and  all  the 
hydrogen  in  the  acid.     The  formulas  of  the  sodium  salts  are  as 
follows:    NaH2PO4,  Na2HPO4,  and  Na3PO4;    the  first  is  called 
monosodium  phosphate  or  primary  sodium  phosphate,  the  second 
disodium  phosphate  or  secondary  sodium  phosphate,  and  the  third 
trisodium  phosphate  or  tertiary  sodium  phosphate. 


PHOSPHORUS,  ARSENIC,  ANTIMONY,  AND  BISMUTH      403 

461.  The  primary  salt  shows  a  slight  acid  reaction;  the  ions 
chiefly  formed  from  the  salt  are  Na+  and  H2PO4~,  but  the  latter 
ionizes  slightly  to  form  H+  and  HPO4~~,  and,  as  a  result,  the  solu- 
tion shows  an  acid  reaction.  The  secondary  salt  shows  a  very 
weak  alkaline  reaction  as  the  result  of  hydrolysis.  In  this  case 
the  ions  formed  in  largest  amount  are  2Na+  and  HPO4~~,  but 
the  latte'r  reacts  to  some  extent  with  the  hydrogen  ions 
formed  from  water,  H20  <=±  H+  -f  OH~,  and  in  uniting  with 
them  to  form  H2PO4~(HPO4~  +  H+<=»H2PO4~)  leaves  an  ex- 
cess of  OH~  ions  in  the  solution,  which,  as  a  result,  shows  an  alka- 
line reaction.  The  tertiary  salt,  which  can  be  made  by  evaporating 
to  dryness  a  solution  of  the  secondary  salt  and  sodium  hydrox- 
ide, is  converted  by  water  into  the  secondary  salt  and  sodium 
hydroxide  : 

Na3PO4  +  H2O  <=>  Na2HPO4  +  NaOH 

The  equation  written  with  ionic  formulas  to  emphasize  the  hydroly- 
sis is  as  follows  : 


3Na+  +  PO4-     '  +  H     +  OH~  <=»3Na     +  HPO4~     +  OH~ 

462.  The  action  of  heat  on  primary  and  secondary  phosphates 
is  represented  by  the  following  equations  : 

NaH2PO4  =  H2O  +  NaPO3 
2Na2HPO4  =  H2O  +  Na4P2O7 

In  both  cases  the  elements  of  water  are  lost  ;  primary  orthophos- 
phates  are  thus  converted  into  metaphosphates,  and  the  secondary 
salts  into  pyrophosphates.  The  reactions  take  place  very  slowly 
in  the  reverse  direction  in  aqueous  solutions  of  metaphosphates  and 
pyrophosphates. 

When    sodium-ammonium    phosphate    (microcosmic    salt)    is 
heated,  both  water  and  ammonia  are  lost: 

NaNH4HPO4  =  NH3  +  H2O  +  NaPO3 

Magnesium-ammonium  phosphate  loses  ammonia  under  the 
same  conditions: 

2MgNH4PO4  =  2NH3  +  H20  +  Mg2P207 


404  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  first-named  salt  is  used  in  qualitative  analysis  (see  borax- 
beads),  and  the  latter  in  quantitative  analysis.  Magnesium- 
ammonium  phosphate,  which  is  insoluble  in  water,  is  the  salt 
commonly  used  to  separate  magnesium  from  other  metals  in 
qualitative  and  quantitative  analysis.  When  it  is  used  for  the 
latter  purpose  it  is  heated  and  weighed  as  the  pyrophosphate. 

The  normal  phosphates,  except  those  of  the  alkali  metals,  are 
insoluble  in  water,  but  dissolve  in  strong  acids;  the  acid  phosphates 
of  the  alkaline  earths  (calcium,  strontium,  etc.)  are  also  soluble  in 
water. 

463.  Tests  for  Orthophosphates. — Orthophosphoric  acid  and  its 
soluble  salts  are  converted  in  solution  by  silver  nitrate  into  silver 
orthophosphate,  Ag3PO4,  which  is  formed  as  a  yellow  precipitate. 
This  test  serves  to  distinguish  the  acid  and  its  salts  from  the  other 
phosphoric   acids   or  their  salts,   which   give  white   precipitates 
under  the  same  conditions;    it  is  not,  however,  characteristic  of 
orthophosphates.     Another    test,    frequently    used,    consists    in 
adding  to  the  solution  of  the  acid  or  salt  a  solution  of  magnesium 
chloride    and    ammonia;     magnesium-ammonium    phosphate    is 
formed  as  a  crystalline  precipitate. 

A  third  test  commonly  applied  in  qualitative  and  quantitative 
analysis  is  based  on  the  formation  of  a  copious  yellow  precipitate 
when  a  solution  of  a  phosphate  is  treated  with  a  solution  of  ammo- 
nium molybdate,  (NH^MoQ*,  containing  nitric  acid.  The 
precipitate  has  the  complex  composition  represented  by  the 
formula  (NH4)3PO4,12MoO3,6H2O,  and  is  called  ammonium  phos- 
phomolybdate.  Orthoarsenic  acid  and  its  salts,  which  resemble 
closely  the  analogous  compounds  of  phosphorus,  give  a  precipitate 
similar  in  appearance  and  composition  to  that  derived  from  ortho- 
phosphoric  acid;  for  this  reason  positive  tests  for  orthophosphates 
should  be  followed  by  a  test  for  arsenic.  Silver  or thoar senate  is  a 
brown  precipitate. 

464.  Phosphorous  Acid,  H3PO3. — The  formation  of  this  acid 
from  phosphorus  trioxide  has  already  been  noted.     It  is  best  pre- 
pared by  the  action  of  phosphorus  trichloride  on  water: 

PC13  +  3H20  =  H3P03  +  3HC1 

Phosphorous  acid  is  a  comparatively  weak  acid  and  forms  salts 
as  the  result  of  the  replacement  of  one  or  two  hydrogen  atoms  only. 


PHOSPHORUS,  ARSENIC,  ANTIMONY,  AND  BISMUTH      405 

It  is  an  active  reducing  agent  and  precipitates  the  noble  metals 
from  solutions  of  their  salts.  When  heated  it  decomposes,  and 
metaphosphoric  acid  and  phosphine,  PH3,  are  formed,  the  change 
in  valence  being  from  +3  in  phosphorous  acid  to  +5  in  meta- 
phosphoric acid  and  —3  in  phosphine. 

465.  Hypophosphorous  Acid,  H3PO2. — Salts  of  hypophosphor- 
ous  acid  are  formed  when  phosphorus  is  dissolved  in  a  solution  of  a 
strong  base.     It  will  be  recalled  that  when  chlorine  reacts  with 
sodium  hydroxide,  sodium  hypochlorite  and- sodium  chloride  are 
formed,  and  the  reaction  was  explained  on  the  assumption  that 
the  halogen  reacted  with  water  and  formed  hypochlorous  acid  and 
hydrochloric  acid,  which  were  subsequently  neutralized  by  the 
base  present  (432).     It  is  probable  that  a  similar  reaction  takes 
place  in  the  case  of  phosphorus,  as  a  result  of  which  hypophos- 
phorous  acid  and  phosphine  are  produced: 

4P  +  6H2O  =  PH3  +  3H3PO2 

Phosphine  is  not  an  acid  and  is  not  affected,  therefore,  by  the 
base  present  and  escapes  as  a  gas;  the  acid  produced,  which  is 
monobasic,  is  converted  into  a  salt.  The  equation  for  the  reaction 
when  white  phosphorus  is  dissolved  in  a  warm  solution  of  sodium 
hydroxide  is,  therefore,  as  follows: 

4P  +  3H2O  +  3NaOH  =  PH3  +  3NaH2PO2 

In  preparing  the  free  acid,  barium  hydroxide  is  used  instead  of 
sodium  hydroxide,  and  the  barium  salt  formed  is  converted  into  the 
acid  and  barium  sulphate  by  adding  sulphuric  acid  to  the  solution; 
on  removing  the  precipitate  by  filtration  a  solution  of  hypophos- 
phorous  acid  is  obtained.  The  acid  is  a  white  crystalline  solid; 
it  is  a  very  active  reducing  agent. 

466.  Phosphine. — -A  reaction  by  which  phosphine  can  be  pre- 
pared from  white  phosphorus  and  a  solution  of  a  base  has  just  been 
discussed    (465)    in    some    detail.     Certain    side-reactions    occur 
which  lead  to  the  formation  of  hydrogen  and  a  second  hydride  of 
phosphorus  of  the  formula  P2H4,  which  is  a  liquid  that  boils  at  57°. 
Since  the  latter  is  spontaneously  inflammable,  the  gas  generated 
takes  fire  when  it  comes  into  contact  with  the  air,  unless  it  has  been 
freed  from  the  liquid  hydride  by  passing  it  through  alcohol  or 
some  other  solvent. 


406  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Phosphine  can  be  prepared  by  decomposing  with  water  the 
phosphides  of  the  more  active  metals;  calcium  phosphide  can  be 
used  conveniently  for  this  purpose : 

Ca3P2  +  6H2O  =  3Ca(OH)2  +  2PH3 

Phosphine  is  a  colorless;  inflammable  gas.  The  products  of  its 
combustion  are  water  and  metaphosphoric  acid,  which  is  produced 
as  the  result  of  the  union  of  the  phosphorus  pentoxide  formed  with 
a  part  of  the  water  produced.  Phosphine  resembles  ammonia  in 
composition  and  like  ammonia  unites  with  the  halogen  hydrides 
to  form  compounds  analogous  in  composition  to  ammonium  salts; 
the  most  stable  of  these  is  phosphonium  iodide,  PELiI.  The 
phosphonium  compounds  are  decomposed  by  water  into  phosphine 
and  the  free  acid.  Phosphine  is  insoluble  in  water  and  does  not 
react  with  it  to  form  a  compound  analogous  to  ammonium  hy- 
droxide. 

Phosphine  resembles,  in  some  respects,  hydrogen  sulphide; 
when  passed  into  solutions  of  the  salts  of  some  of  the  metals  it 
causes  the  precipitation  of  insoluble  phosphides;  the  case  of 
copper  is  an  example : 

3CuSO4  +  2PH3  =  Cu3P2  +  3H2SO4 

467.  Halides  of  Phosphorus. — Phosphorus  trichloride,  PC13,  is 
formed  when  chlorine  is  passed  over  white  phosphorus,  care  being 
taken  to  avoid  an  excess  of  the  halogen;  under  the  latter  circum- 
stances the  pentachloride,  PCls,  is  formed,  since  it  is  produced  as 
the  result  of  the  direct  action  of  the  trichloride  and  chlorine.  The 
trichloride  is  a  heavy,  colorless  liquid,  which  boils  at  76°.  It 
reacts  with  water  to  form  phosphorous  acid  and  hydrochloric  acid, 
and  in  the  hydrolysis  resembles  the  halides  of  other  acid-forming 
elements. 

Phosphorus  pentachloride  is  a  white  solid,  which  sublimes  at 
163°,  and  dissociates  into  the  trichloride  and  chlorine  at  higher 
temperatures.  It  is  hydrolyzed  very  rapidly  by  water,  the  product 
formed  being  determined  by  the  amount  of  water  present.  With 
excess  of  water  complete  hydrolysis  takes  place  and  phosphoric 
acid  and  hydrochloric  acid  are  formed : 

PC15  +  4H2O  =  H3PO4  -f  5HC1 


PHOSPHORUS,  ARSENIC,  ANTIMONY,  AND  BISMUTH      407 

If  not  enough  water  is  present  to  bring  about  this  reaction  partial 
hydrolysis  takes  place  and  phosphorus  oxychloride  is  formed: 

PC15  +  H2O  =  POC13  +  2HC1 

Phosphorus  forms  with  fluorine  and  with  bromine  compounds  in 
which  it  shows  the  valencies  3  and  5.  The  two  iodides  have  the 
formulas  P2l4  and  Pis.  All  these  compounds  resemble  the  chlo- 
rides in  chemical  behavior  and  are  prepared  by  the  same  general 
methods  as  those  used  to  prepare  the  latter. 

468.  Sulphides  of  Phosphorus. — Compounds  represented  by 
the  following  formulas  can  be  prepared  by  heating  phosphorus 
with  the  proper  amounts  of  sulphur:  P4S3,  P4S7,  and  P2Ss     They 
are  all  solids  which  are  hydrolyzed  by  water.     Phosphorus  pen- 
tasulphide  reacts  with  the  latter  and  forms  phosphoric  acid  and 
hydrogen  sulphide: 

P2S5  +  8H2O  =  2H3PO4  +  5H2S 
The  sulphide  of  the  formula  P4S3  is  used  in  making  matches. 

ARSENIC 

469.  The  chemistry  of  arsenic  and  its  compounds  resembles 
very  closely  that  of  phosphorus  and  its  compounds.     Since  the 
latter  has  just  been  discussed  at  length  it  will  be  possible  to  limit 
largely  the  consideration  of  arsenic  to  an  account  of  any  unique 
properties  exhibited  by  it  or  its  compounds,  and  to  the  uses  to 
which  they  are  put. 

Arsenic  is  a  much  less  active  element  than  phosphorus  and 
is  less  strongly  electro-negative  in  character.  Its  compounds  with 
hydrogen  and  with  oxygen  are  less  stable  than  the  corresponding 
phosphorus  derivatives. 

470.  History  and  Occurrence. — Two  sulphides  of  arsenic  occur 
as  minerals  which  are  highly  colored  and,  therefore,  early  attracted 
attention;   orpiment,  As2S3,  is  bright  yellow,  and  realgar,  As2S2, 
is  orange-red.     When  heated  in  the  air  both  compounds  are  con- 
verted into  arsenic  trioxide,  As2O3,  which  occurs  also  as  a  mineral 
called  arsenite.     The  properties  of  these  substances  were  known 
to  the  early  Greeks,  who  gave  to  orpiment  a  name  from  which  the 
word  arsenic  is  derived.     The  element  in  the  free  condition  was 


408  INORGANIC  CHEMISTRY  FOR  COLLEGES 

known  to  the  alchemists  in  the  thirteenth  century,  a  fact  which 
resulted,  no  doubt,  from  the  ease  with  which  arsenic  oxide  is 
reduced  when  heated  with  charcoal.  The  poisonous  properties  of 
the  oxide  have  been  known  and  used  since  the  earliest  times. 

Arsenic  occurs  free  in  nature,  in  combination  with  metals  as 
arsenides  such  as  smaltite,  CoAs2,  and  in  association  with  sulphur 
in  metallic  sulphides.  Arsenical  pyrite,  mispickel,  FeAsS,  is  an 
important  mineral,  which  occurs  along  with  pyrite,  FeS2;  zinc 
sulphide,  sphalerite,  ZnS,  also  contains  compounds  in  which  the 
sulphur  is  in  part  replaced  by  arsenic;  cobaltite  has  the  formula 
CoAsS.  When  pyrite  is  burned  to  make  sulphur  dioxide  in  the 
manufacture  of  sulphuric  acid,  the  arsenic  in  it  is  converted  into 
arsenic  trioxide,  and  since  the  latter  is  a  solid  it  condenses  in  this 
form  in  the  dust  boxes  and  flues.  In  the  smelting  of  zinc  ores 
containing  arsenic,  the  latter  is  obtained  as  the  trioxide  when  the 
ores  are  heated  in  the  air  (roasted)  preparatory  to  reduction.  The 
oxide  obtained  from  these  sources  is  heated  with  carbon  and  the 
element  liberated  is  distilled  off  and  condensed.  Arsenic  is  also 
obtained  by  heating  arsenical  pyrites;  arsenic  distills  off  and  fer- 
rous sulphide  is  left  in  the  retort. 

471.  Physical  Properties. — Arsenic  has  the  appearance  of  a 
metal;   it  is  steel-gray  in  color,  is  more  or  less  brittle,  and  when 
fractured,  shiny  crystalline  surfaces  can  be  seen.     It  has  the  den- 
sity 5.7,  and  is  a  poor  conductor  of  electricity.     Arsenic  sublimes 
at  relatively  low  temperatures,  and  at  600°  the  pressure  of  its  vapor 
is  equal  to  1  atmosphere.     The  vapor  has  a  light-yellow  color 
and  its  density  indicates  that  its  formula  is  As4;    when  cooled 
quickly  it  condenses  to  a  yellow  modification  of  the  element,  which 
resembles   closely   white   phosphorus  in  properties,  being  phos- 
phorescent and  soluble  in  carbon  disulphide. 

472.  Chemical  Properties. — When  arsenic  burns  in  air  or  oxy- 
gen the  trioxide,  As2Os,  is  formed.     It  unites  with  the  halogens 
and  sulphur  readily  and  with  the  metals  forms  arsenides.     Active 
oxidizing  agents  convert  it  into  arsenic  acid,  HaAsCU.     The  ele- 
ment does  not  displace  hydrogen  from  acids. 

473.  Oxides  of  Arsenic. — Arsenic  trioxide,   As2C>3,   which  is 
formed  when  arsenic  burns  or  when  arsenides  are  heated  in  the  air, 
has  been  known  for  a  long  time  and  is  commonly  called  "  arsenic  " 
or  white  arsenic.     It  is  purified  by  sublimation  and  is  obtained 


PHOSPHORUS,  ARSENIC,  ANTIMONY.  AND  BISMUTH      409 

as  a  white  crystalline  powder.  If  the  vapor  is  cooled  rapidly  a 
transparent  form  resembling  glass  is  produced,  which  passes  slowly 
into  the  crystalline  variety.  The  oxide  dissolves  to  a  slight  degree 
in  hot  water  and  reacts  with  the  latter  to  form  arsenious  acid, 
HaAsOa.  It  is  more  soluble  in  strong  hydrochloric  acid  with 
which  it  forms  arsenic  trichloride;  with  concentrated  sulphuric 
acid,  a  basic  sulphate  is  formed,  which  is  hydrolyzed  by  water. 

Arsenic  trioxide  is  an  active  poison;  the  fatal  dose  is  from  0.06 
to  0.18  gram  (1  to  3  grains).  Many  people  who  live  in  high  alti- 
tudes are  said  to  eat  white  arsenic  as  it  is  a  help  in  respiration, 
which  is  more  or  less  difficult  in  places  where  the  concentration  of 
the  oxygen  in  the  air  is  low.  A  solution  of  the  oxide  in  sodium 
carbonate  (Fowler's  Solution)  is  administered  as  a  stimulant  in 
certain  nervous  affections. 

Arsenic  trioxide  is  used  in  making  high-quality  colorless  glass; 
it  serves  as  an  oxidizing  agent,  the  arsenic  produced  being  vola- 
tilized. It  is  used  to  some  extent  as  a  mordant  in  calico-printing, 
in  making  pigments,  as  an  ingredient  of  poisons  for  rats  and  flies 
and  other  insects,  and  as  a  preservative  for  untanned  hides. 

Arsenic  pentoxide  is  obtained  as  a  white  crystalline  substance 
by  cautiously  heating  arsenic  acid.  It  dissolves  in  water  and 
reacts  with  it  to  form  arsenic  acid. 

474.  Acids  of  Arsenic. — Salts  containing  arsenic  are  known 
which  resemble  in  composition  and  physical  and  chemical  proper- 
ties the  salts  of  ortho-,  pyro-,  and  meta-phosphoric  acids  and 
phosphorous   acid.      Orthoarsenic   acid,    (H3AsO4)2,H2O,   is   the 
only  acid  of  arsenic  which  has  been  isolated.     Silver  orthoarsenate, 
which  is  obtained  as  a  brown  precipitate,  and  magnesium-ammo- 
nium arsenate  are  characteristic  salts  that  resemble  the  corre- 
sponding  salts   of   orthophosphoric    acid.     Arsenious   acid    is    a 
very  weak  acid  and  its  soluble  salts,  which  are  formed  by  dissolving 
arsenic  trioxide  in  solutions  of  the  caustic  alkalies,  are  highly 
hydrolyzed ;  the  arsenites  of  the  heavy  metals  prepared  from  these 
by  double  decomposition  possess,  in  most  cases,   complex  for- 
mulas. 

475.  Antidotes  for  Poisons. — When  a  poison  has  been  taken 
into  the  body  through  the  mouth,  the  usual  procedure  is  to  admin- 
ister as  soon  as  possible  an  emetic,  which  causes  vomiting.     An 
antidote  is  next  taken,  the  object  of  which  is  to  convert  the  poison 


410  INORGANIC  CHEMISTRY  FOR  COLLEGES 

into  an  insoluble  substance  and  thus  reduce  as  much  as  possible 
the  absorption  of  the  poison  by  the  tissues  of  the  body.  In  the 
case  of  white  arsenic  or  arsenites,  freshly  precipitated  ferric  hydrox- 
ide or  magnesium  hydroxide  is  commonly  used  on  account  of  the 
fact  that  ferric  arsenite  and  magnesium  arsenite  are  insoluble 
salts.  The  frequent  use  of  white  of  egg  as  an  antidote  is  based  on 
the  fact  that  the  salts  of  the  heavy  metals  which  are  poisonous 
form,  in  most  cases,  insoluble  precipitates  with  the  proteins 
present  in  the  egg. 

476.  Arsine. — The  arsenides  of  the  more  active  metals  react 
with  hydrochloric  acid  and  form  chlorides  and  arsine,  AsHa. 
Sodium  arsenide  and  magnesium  arsenide  are  hydrolyzed  by 
water  and,  thus,  resemble  the  analogous  nitrides  and  phosphides. 
Arsine  can  be  prepared  in  these  ways.  It  is  a  gas  with  a  dis- 
agreeable garlic-like  odor,  and  is  very  poisonous.  It  liquefies 
at  -40°. 

Arsine  is  formed  when  compounds  containing  arsenic  are 
reduced  by  nascent  hydrogen,  and  it  is  this  method  which  is  used 
in  testing  for  arsenic,  especially  when  it  is  present  in  small 
quantities  in  the  substance  under  examination.  Arsine  is  readily 
decomposed  by  heat  into  the  elements  of  which  it  is  composed, 
and  when  the  decomposition  is  effected  in  a  glass  tube  arsenic  is 
deposited  in  the  form  of  a  brown  coating,  or  as  a  mirror  if  enough 
of  the  element  is  present.  This  method  of  detecting  arsenic  is 
known  as  Marsh's  test,  and  is  the  one  commonly  used  in  col- 
lecting evidence  to  be  presented  in  court  in  cases  of  suspected 
poisoning  by  arsenic.  In  carrying  out  the  test,  pure  zinc  and  pure 
dilute  sulphuric  acid  are  placed  in  a  hydrogen  generator,  and  the 
gas  formed  passed  first  through  a  tube  containing  dehydrated  cal- 
cium chloride,  which  is  used  as  a  drying  agent,  and  then  through  a 
tube  of  hard  glass  which  is  constricted  at  one  place.  When  hydro- 
gen is  freely  evolved  the  hard  glass  tube  is  heated  with  a  Bunsen 
flame  at  a  point  between  the  generator  and  the  constricted  part 
of  the  tube.  If,  after  a  few  minutes,  no  deposit  is  formed  in  the 
hot  tube,  the  solution  to  be  tested  for  arsenic  is  poured  through 
the  thistle  tube  into  the  generator.  The  preliminary  heating  is 
necessary,  because  most  samples  of  zinc  and  of  sulphuric  acid  con- 
tain traces  of  arsenic,  which  are  present  as  the  result  of  the  fact 
that  both  of  these  substances  are  made  from  raw  materials  that 


PHOSPHORUS,  ARSENIC,  ANTIMONY,  AND  BISMUTH      411 

are  apt  to  contain  arsenic.  If  the  compound  introduced  into  the 
generator  contains  arsenic,  it  is  reduced  by  the  nascent  hydrogen; 
the  arsine  formed  is  decomposed  in  the  hot  part  of  the  tube  and 
arsenic  is  deposited  in  the  constriction  of  the  latter.  In  using 
this  method  to  determine  the  arsenic  quantitatively,  a  series 
of  tubes  is  prepared  in  which  arsenic  is  deposited  from  known 
amounts  of  a  solution  of  sodium  arsenite.  By  comparing  the  size 
and  appearance  of  the  deposit  on  these  standards  with  the  tube 
used  in  the  case  of  the  substance  being  tested,  it  is  possible  to 
determine  the  quantity  of  arsenic  present. 

When  compounds  containing  antimony  are  heated  with  nas- 
cent hydrogen  a  volatile  hydride,  easily  decomposed  by  heat,  is 
formed.  It  is  possible,  however,  to  distinguish  the  deposit  obtained 
in  this  case  from  that  produced  by  arsenic  (487). 

Arsine  burns  when  ignited,  and  is  converted  into  water  and 
arsenic  trioxide,  which  forms  a  white  cloud.  If  a  piece  of  porce- 
lain is  put  into  an  arsine  flame,  a  black  deposit  of  arsenic  is  formed 
on  the  cold  surface,  the  phenomenon  observed  being  similar  to 
that  which  occurs  when  a  cold  object  is  brought  into  contact  with 
an  illuminating  gas  flame  (195). 

Arsine  is  insoluble  in  water,  does  not  react  with  aqueous  solu- 
tions of  acids  to  form  salts  as  ammonia  does,  and  does  not  exhibit 
the  behavior  shown  by  phosphine  in  combining  with  the  halogen 
hydrides. 

477.  Halides  of  Arsenic. — The  compounds  of  arsenic  and  the 
halogens  have  the  general  formula  AsXs.     They  are  hydrolyzed 
by  water,  but  owing  to  the  slight  basic  properties  of  arsenious  acid 
the  reactions  are  reversible.     Arsenic  trichloride  boils  at   130°, 
and  when  a  solution  of  it  in  hydrochloric  acid  is  distilled,  it  passes 
over  with  the  acid  and  the  water;   it  is  thus  possible  to  separate 
arsenic  in  this  way  from  many  other  elements. 

478.  Sulphides  of  Arsenic. — The  occurrence  of  two  sulphides 
of  arsenic  as  minerals  has  already  been  noted.     Artificial  orpiment, 
As2S3,  is  made  by  subliming  a  mixture  of  arsenic  trioxide  and  sul- 
phur and  is  used  as  a  yellow  pigment.     Realgar,  As2S2,  is  made  by 
fusing  together  the  oxide  and  sulphur,  or  by  distilling  arsenical 
ores  with  sulphur.     It  is  used  to  some  extent  as  a  red  pigment,  but 
its  chief  use  is  in  making  mixtures  to  produce  colored  lights  when 
burned,  and  for  removing  the  hair  from  hides  preparatory  to  tanning. 


412  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Arsenic  trisulphide  is  formed  as  a  yellow  precipitate  when 
hydrogen  sulphide  is  passed  into  a  solution  of  arsenic  trichloride. 
It  is  separated  in  this  way  in  the  procedure  used  in  qualitative 
analysis.  The  sulphide  is  practically  insoluble  in  hot  concentrated 
hydrochloric  acid,  a  fact  which  is  used  in  separating  it  from  anti- 
mony trisulphide,  which  dissolves  in  this  reagent. 

Arsenic  trisulphide  dissolves  in  ammonium  sulphide  as  the 
result  of  the  formation  by  direct  addition  of  a  compound,  which  is 
soluble  in  water: 

As2S3  +  3(NH4)2S  =  As2S3,3(NH4)2S 

The  formula  of  the  compound  is  generally  written  as  (NH4)3AsS3, 
and  it  is  called  ammonium  thioarsenite;  it  can  be  considered 
as  the  ammonium  salt  of  the  acid  H3AsS3  formed  as  the  result  of 
the  replacement  of  the  oxygen  in  arsenious  acid,  H3AsO3,  by  sul- 
phur. Many  salts  of  such  thioacids  are  known.  When  an 
acid  is  added  to  the  solution  of  ammonium  thioarsenite  the  salt 
is  decomposed: 

(NH4)3AsS3  +  3HC1  =  3NH4C1  +  H3AsS3 
2H3AsS3  =  As2S3  +  3H2S 

Arsenic  trisulphide  is  precipitated  and  hydrogen  sulphide  is  set  free. 
479.  When  arsenic  trisulphide  is  treated  with  a  solution  of 
yellow  ammonium  sulphide,  which  contains  the  polysulphide  (284), 
some  of  the  sulphur  of  the  latter  unites  with  the  trisulphide  to 
form  the  pentasulphide  of  arsenic,  which,  in  turn,  reacts  with 
ammonium  sulphide  to  form  the  compound  As2Sr,,3(NH4)2S. 
The  formula  of  this  compound  is  written  (NH4)3AsS4  and  it  is 
called  ammonium  thioarsenate;  it  bears  the  same  relation  to 
ammonium  arsenate,  (NH4)3AsO4,  that  ammonium  thioarsenite, 
(NH4)3AsS3,  bears  to  ammonium  arsenite,  (NH4)3AsO3.  When 
a  solution  of  the  thioarsenate  is  treated  with  an  acid,  arsenic 
pentasulphide  is  precipitated : 

(NH4)3AsS4  +  3HC1  =  3NH4C1  +  H3AsS4 
2H3AsS4  =  As2S5  +  3H2S 

These  reactions  are  used  in  qualitative  analysis.     The  salts  are 
often  called  sulpharsenites  and  sulpharsenates. 


PHOSPHORUS,  ARSENIC,  ANTIMONY,  AND  BISMUTH      413 

When  hydrogen  sulphide  is  passed  into  a  solution  of  arsenic 
acid  in  concentrated  hydrochloric  acid,  arsenic  pentasulphide  is 
precipitated. 

ANTIMONY 

480.  A  comparison  of  the  chemistry  of  antimony  with  that  of 
phosphorus  and  arsenic  brings  out  clearly  the  progressive  change  in 
properties  within  a  chemical  family,  with  increase  in  atomic  weight. 
Phosphorus  acts  as  an  acid-forming  element  only.  Arsenious 
hydroxide  is  both  a  base  and  an  acid.  When  the  hydroxide  of  an 
element  shows  both  acidic  and  basic  properties  the  element  is 
said  to  be  amphoteric.  We  shall  see  that  antimonous  acid  is 
a  weaker  acid  than  arsenious  acid,  but  that  it  is  a  stronger  base 
than  the  latter,  since  its  salts  cannot  be  completely  hydrolyzed  by 
cold  water.  The  pentoxide  of  antimony  acts  as  an  acid  anhy- 
dride only.  This  fact  serves  as  an  example  of  the  generalization 
that  an  increase  in  the  valence  of  an  element  toward  oxygen  is  asso- 
ciated with  an  increase  in  acidic  properties,  and,  as  a  result,  a 
decrease  in  basic  properties.  Of  two  acids  derived  from  the  same 
element  the  stronger  acid  is  the  one  in  which  the  element  shows  the 
higher  valence. 

There  is  a  gradation  in  the  properties  of  the  hydrides  of  the 
elements  in  the  phosphorus  family.  Since  hydrogen  acts  as  a  posi- 
tive element,  the  most  stable  hydrides  are  those  in  which  it  is  com- 
bined with  strongly  negative  elements.  As  the  negative  character 
of  the  elements  in  the  phosphorus  family  drops  off,  the  stability  of 
the  hydrides  decreases;  antimony  hydride  is  decomposed  into  its 
elements  at  a  much  lower  temperature  than  is  arsine.  Similar 
relationships  exist  among  other  compounds  of  these  elements, 
which  it  will  be  of  interest  to  the  student  to  discover. 

481.  History. — Compounds  of  antimony  have  been  known 
since  Biblical  times,  and  since  they  were  used  for  a  number  of  pur- 
poses they  were  studied  in  considerable  detail  by  the  alchemists 
and  by  their  successors,  the  so-called  iatrochemists,  who  centered 
their  attention  on  the  use  of  chemical  substances  as  drugs.  The 
effect  on  the  body  of  many  inorganic  compounds  and  organic 
substances  isolated  from  plants  was  studied  by  the  iatrochemists 
in  the  Middle  Ages,  and  the  facts  discovered  became  the  basis 
of  the  branch  of  medicine  called  materia  medica,  which  has  to  do 


414  INORGANIC  CHEMISTRY  FOR  COLLEGES 

with  the  materials  used  in  the  science  on  account  of  their  effect 
on  the  body.  Many  of  the  substances  prescribed  to-day  as  drugs 
were  first  studied  by  the  iatrochemists,  of  whom  Paracelsus  was  the 
leader.  Basil  Valentine  first  described  in  the  fifteenth  century  a 
method  of  preparing  antimony. 

482.  Occurrence. — Antimony  occurs  free  in  small  quantities, 
usually   in   association   with   arsenic.     Its   chief   ore   is   stibnite, 
Sb2Sa,  which  is  found  in  the  form  of  large  glistening  black  prisms. 
Two  oxides,  Sb2Os  and  SboCU,  occur  as  minerals.     The  element, 
like  arsenic,  is  associated  with  sulphur  in  certain  metallic  sulphides, 
and  is  obtained  as  a  by-product  in  the  metallurgy  of  these  metals, 
of  which  lead  and  copper  are  examples. 

483.  Preparation. — Antimony  is  obtained  from  stibnite  either 
by  direct  reduction  with  iron,  or  by  first  roasting  the  mineral  to 
convert  it  into  the  oxide,  and  then  reducing  the  latter  with  carbon. 
In  the  first  process  the  ore  is  mixed  with  salt  and  scrap  iron  in  a 
large  crucible  and  heated  to  a  high  temperature  in  a  furnace.     The 
salt  melts  and  serves  as  a  flux  (521),  and  the  heavy  ferrous 
sulphide  sinks  to  the  bottom  of  the  crucible.     The  reaction  which 
takes  place  is  represented  by  the  following  equation : 

Sb2S3  +  3Fe  =  2Sb  +  3FeS 

The  antimony  formed  is  poured  off  and  cast  into  a  convenient 
form.  Since  it  contains  several  per  cent  of  iron,  it  must  be  purified ; 
this  is  done  by  melting  it  with  just  enough  stibnite  to  react  with  the 
iron  present. 

484.  Properties. — Antimony  is  a  bluish-white,  highly  crystal- 
line substance  with  a  marked  luster.     It  can  be  readily  powdered, 
and  has  a  high  density,  6.7;   it  melts  at  630°  and  boils  at  1440°. 
Its  vapor  just  above  its  boiling-point  is  made  up  of  molecules 
containing  3  and  4  atoms  of  the  element.     The  chief  use  of  anti- 
mony is  in  the  preparation  of  alloys  such  as  type-metal  and  brit- 
tania  metal,  which  are  to  be  cast  in  molds.     Owing  to  the  presence 
of  antimony  these  alloys  expand  on  solidification  and  give  sharp 
castings. 

Antimony  burns  in  the  air  and  forms  the  trioxide,  and  with 
chlorine  and  bromine  it  forms  trihalides.  It  combines  directly  with 
sulphur,  arsenic,  phosphorus,  and  some  metals.  It  is  oxidized  by 
concentrated  nitric  acid  to  the  trioxide,  which  on  continued  heating 


PHOSPHORUS,  ARSENIC,  ANTIMONY,  AND  BISMUTH      415 

with  the  acid  is  changed  to  antimonic  acid.     It  is  converted  by 
hot  concentrated  sulphuric  acid  into  antimony  sulphate,  802(804)3. 

485.  Oxides    and    Salts    of    Antimony. — When    antimony    is 
burned  in  the  air  the  chief  product  is  the  trioxide,  SboOa,  which  is 
white;  if  an  excess  of  oxygen  is  used  the  tetroxide,  Sb2O4,  is  formed. 
The  latter  possesses  neither  acidic  nor  basic  properties.     Anti- 
mony pentoxide,  Sb2O5,  is  obtained  as  a  yellow  powder  when  anti- 
monic acid  is  heated;   it  dissolves  in  bases  and  forms  antimonates. 

Antimony  trioxide  dissolves  in  alkalies  and  antimonites  are 
formed.  It  reacts  with  acids  to  form  salts;  it  is  converted,  for 
example,  by  strong  hydrochloric  acid  into  antimony  trichloride, 
SbCls,  by  nitric  acid  into  antimony  nitrate,  Sb(NO3)3,  and  by 
sulphuric  acid  into  the  sulphate,  Sb2 (804)3. 

In  all  these  reactions  the  oxide  shows  basic  properties  and  anti- 
mony acts  as  a  metallic  element.  It  is,  however,  weakly  basic 
only,  and,  as  a  consequence,  the  salts  are  hydrolyzed  by  water. 
When  water  is  added  to  a  strong  solution  of  the  chloride,  antimony 
oxy chloride  is  precipitated: 

SbCl3  +  H2O  <=>  SbOCl  +  2HC1 

The  reaction  is  a  reversible  one  and  the  concentrations  at  equilib- 
rium are  determined  by  the  amount  of  acid  present  in  the  solution. 
A  large  number  of  basic  salts  of  antimony  are  known  which 
contain  the  element  and  oxygen  in  combination  as  the  group  SbO. 
For  this  reason  the  latter  has  been  given  a  special  name  and  is 
called  antimonyl;  it  may  be  considered  as  having  the  valence  1. 
Tartar-emetic,  which  has  been  used  in  medicine  for  a  long  time,  is 
prepared  by  treating  cream  of  tartar,  acid  potassium  tartrate, 
KH^EUOe),  with  antimony  trioxide;  the  acidic  hydrogen 
is  replaced  by  the  univalent  SbO  group,  and  the  salt  formed, 
KSbO^HUOe),  ^fbO,  is  potassium  antimonyl  tartrate.  Tartar- 
emetic  is  extensively  used  as  a  mordant  in  dyeing  vegetable  fabrics, 
such  as  cotton  and  linen. 

486.  Acids  of  Antimony! — When  bases  are  added  to  solutions 
of    antimony    salts,    a    hydroxide,   which    is    antimonous    acid, 
Sb(OH)3,  is  precipitated.     The  acid  soon  loses  water  and  passes 
into  the  trioxide. 

Antimonic  acid  is  formed  by  oxidation  of  antimony  with  nitric 
acid,  or  as  the  result  of  the  hydrolysis  of  antimony  pentachloride, 


416  INORGANIC  CHEMISTRY  FOR  COLLEGES 


It  is  a  white  insoluble  compound,  which  dissolves  in  solu- 
tions of  the  alkalies.  The  best  known  salts  of  antimonic  acid  are 
derived  from  the  meta  and  pyro  acid.  Sodium  pyroantimonate, 
Na2HoSb2O7,  is  difficultly  soluble  in  water:  this  fact  is  noteworthy, 
as  most  sodium  salts  are  readily  soluble  in  water. 

487.  Stibine.  —  The  methods  of  preparation  and  properties  of 
stibine,  SbHs,  resemble  closely  those  of  arsine.     The  deposit  of 
the  element  obtained  in  the  Marsh  test,  or  when  a  cold  object  is 
brought  into  contact  with  a  flame  of  burning  stibine,  differs  in 
appearance  and  properties  from  that  produced  by  arsenic.     The 
deposit  in  the  case  of  arsenic  is  shiny  and  brownish-black;   it 
dissolves  in  a  solution  of  bleaching  powder,  and  is  readily  volatil- 
ized at  the  temperature  of  the  Bunsen  flame.     The  deposit  of 
antimony  is  dull  black,  is  insoluble  in  bleaching  powder,  and  does 
not  volatilize  at  as  low  a  temperature  as  arsenic.     Stibine  boils 
at  -17°  and  freezes  at  -88°. 

488.  Halides  of  Antimony.  —  Antimony  trichloride  was  called 
by  the  alchemists  "  butter  of  antimony  "  on  account  of  the  fact 
that  its  crystals  form  a  pasty  mass  in  the  air.     It  is  hydrolyzed 
by  water  and  yields  several  basic  chlorides,  the  best  characterized 
of  which  has  the  composition  SbOCl. 

Antimony  pentachloride  is  a  heavy,  colorless  liquid  which 
fumes  in  the  air;  it  dissolves  in  water,  from  which  it  can  be  ob- 
tained in  the  form  of  hydrated  crystals.  Compounds  of  anti- 
mony with  bromine,  iodine,  and  fluorine  are  known. 

489.  Sulphides  of  Antimony.  —  The  trisulphide  occurs  in  nature 
as  stibnite,  Sb2Ss,  which  is  a  black  mineral.     It  is  formed  as  an 
orange-red  precipitate  when  hydrogen  sulphide  is  passed  into  a 
solution  of  an  antimony  salt.     The  product  prepared  in  this  way 
is  used  as  a  pigment  and  for  vulcanization  in  making  the  red 
antimony  rubber  of  commerce. 

Antimony  trisulphide  dissolves  in  ammonium  sulphide  and  in 
ammonium  polysulphide,  the  reactions  being  analogous  to  those 
in  the  case  of  arsenic  trisulphide.  When  the  thioantimonate 
is  treated  with  an  acid,  antimony  pentasulphide,  Sb2Ss,  is  formed 
as  an  orange-red  precipitate;  it  is  also  formed  when  hydrogen  sul- 
phide is  passed  into  a  solution  of  an  antimonate  in  hydrochloric 
acid. 

Antimony  trisulphide  is  soluble  in  concentrated  hydrochloric 


PHORPHORUS,  ARSENIC,  ANTIMONY,  AND  BISMUTH      417 

acid,  whereas  arsenic  trisulphide  is  not;  the  reaction  serves,  thus,  to 
separate  the  two  elements : 

Sb2S3  +  6HC1  +±  2SbCl3  +  3H2S 

Antimony  pentasulphide  is  also  soluble  in  concentrated  hydro- 
chloric acid,  the  trichloride,  sulphur,  and  hydrogen  sulphide  being 
formed. 

BISMUTH 

490.  Bismuth  is  more  metallic  in  character  than  antimony. 
It  does  not  form  a  hydride  and  its  trioxide  is  not  soluble  in  bases 
and  does  not  act  as  an  acid  anhydride.     Bismuth  trioxide  dissolves 
in  acids  to  form  salts,  which  are  more  or  less  hydrolyzed  by  water. 
Bismuth  pentoxide,  Bi2Os,  has  very  weak  acidic  properties. 

491.  Occurrence   and  Preparation. — Bismuth  is  a   compara- 
tively rare  element;  it  occurs  in  the  free  condition  and  as  the  triox- 
ide, Bi2Os,  (bismuth  ocher),  and  as  the  trisulphide,  61283,  (bis- 
muth glance),  associated  with  silver,  cobalt,    nickel,    and  arsenic 
ores.     The  element  has  been  known  for  centuries  and  was  described 
by  Basil  Valentine. 

The  metal  is  usually  obtained  by  first  roasting  the  ore  in  the 
air,  as  the  result  of  which  the  sulphur  and  most  of  the  arsenic 
present  are  converted  into  oxides  and  volatilized.  The  product  is 
next  heated  with  carbon  (coal),  iron,  and  a  flux,  and  the  metallic 
bismuth  formed  by  reduction  is  drawn  off  while  still  in  the  liquid 
condition.  To  remove  the  arsenic  and  antimony  mixed  with  the 
metal,  it  is  fused  with  sodium  carbonate  and  potassium  nitrate; 
the  latter  oxidizes  the  arsenic  and  antimony  and  converts  them 
into  salts,  but  does  not  affect  the  bismuth. 

492.  Properties. — Bismuth  has  a  silvery  luster  with  a  reddish 
tinge;   it  is  very  brittle,  melts  at  269°,  boils  at  1420°,  and  has  the 
density  9.82.     Its  molecular  formula  is  Bi2,  but  at  very  high  tem- 
peratures the  molecule  undergoes  partial  dissociation  into  atoms. 
It  expands  when  passing  from  the  liquid  to  the  solid  condition, 
whereas  the  common  behavior  is  the  reverse  in  the  case  of  metals. 

Bismuth  is  used  in  making  the  so-called  fusible  alloys.  When 
fused  with  other  low-melting  metals,  a  mixture  is  obtained  which 
melts  at  a  much  lower  temperature  than  that  at  which  any  of  its 
constituents  melts.  Wood's  metal,  for  example,  melts  at  60.5°;  it 


418  INORGANIC  CHEMISTRY  FOR  COLLEGES 

is  made  by  melting  together  4  parts  of  bismuth,  2  parts  of  lead,  1 
part  of  tin,  and  1  part  of  cadmium.  It  is  possible  by  varying 
the  composition  of  the  mixtures  to  obtain  products  that  melt  at 
any  desired  temperature.  The  safety  plugs  used  in  boilers,  elec- 
tric fuses,  automatic  sprinklers,  and  fire-doors  are  made  of  alloys 
prepared  in  this  way. 

Bismuth  unites  with  the  halogens  to  form  salts  of  the  general 
formula  BiXs,  and  is  converted  into  salts  when  heated  with  oxi- 
dizing acids. 

493.  Oxides  of  Bismuth. — The  most  important  oxide  is  the 
trioxide,  Bi2Os,  which  can  be  prepared  by  burning  the  element  or 
by  heating  to  a  high  temperature  the  hydroxide  or  nitrate  of  the 
metal;    it  is  a  yellow  powder  which  dissolves  in  strong  acids  to 
form  salts.     When  bismuth  hydroxide,   suspended  in  water,    is 
treated  with  a  solution  of  stannous  chloride,  SnCk,  a  black  pre- 
cipitate is  formed,  to  which  the  formula  BiO  is  assigned.     If  the 
hydroxide  is  treated  with  chlorine-water,  oxides  are  formed  which 
have  the  composition  Bi2O4  and  Bi2Os. 

494.  Salts  of  Bismuth. — Bismuth  chloride,  BiCl3,H2O,  can  be 
prepared  by  dissolving  the  oxide  in  an  excess  of  hydrochloric  acid 
and  evaporating  the  solution  to  crystallization.     Bismuth  nitrate, 
Bi(NOs)3,5H2O,  is  prepared  in  a  similar  way.     Both  salts  are 
hydrolyzed  by  water;    in  the  case  of  the  chloride  two  chlorine 
atoms  are  replaced  by  hydroxyl  groups: 

BiCl3  +  2H2O  ^±  Bi(OH)2Cl  +  2HC1 

A  similar  reaction  takes  place  in  the  case  of  the  nitrate.  The 
basic  chloride  is  converted  into  bismuth  oxychloride,  BiOCl,  when 
dried.  The  basic  nitrate,  Bi(OH)2NO3,  is  used  in  medicine, 
under  the  name  subnitrate  of  bismuth,  as  an  internal  remedy  in 
the  case  of  certain  stomach  and  intestinal  troubles.  It  is  also  used 
in  face-powders. 

Bismuth  sulphide,  Bi2Ss,  is  obtained  as  a  brownish-black  pre- 
cipitate when  hydrogen  sulphide  is  passed  into  a  solution  of  a  bis- 
muth salt;  the  presence  of  a  dilute  acid  does  not  interfere  with  its 
precipitation.  The  sulphide  is  not  soluble  in  ammonium  poly- 
sulphide  and  can  be  separated  by  means  of  this  reagent  from 
arsenic  and  antimony.  The  insolubility  of  the  sulphide  is  due  to 
the  fact  that  bismuth  is  more  metallic  in  character  than  the  other 


PHOSPHORUS,  ARSENIC,  ANTIMONY,  AND  BISMUTH      419 

members  of  the  phosphorus  family  and  does  not  exhibit  the  prop- 
erty of  acting  as  an  acid-forming  element  in  compounds  analogous 
to  the  thioarsenites  and  thioarsenates. 

495.  Tests  for  Bismuth.  —  In  the  usual  procedure  followed  in 
qualitative  analysis,  bismuth  is  precipitated  as  the  sulphide  from 
a  hydrochloric  acid  solution  along  with  the  other  elements 
which  form  sulphides  insoluble  in  acids.  When  the  sulphides  are 
treated  with  warm  ammonium  polysulphide  to  remove  the  sul- 
phides of  arsenic,  antimony,  and  tin,  bismuth  sulphide  is  not 
affected  as  it  is  insoluble  in  this  reagent.  When  the  insoluble 
sulphides  left  after  the  treatment  are  heated  with  nitric  acid,  bis- 
muth sulphide  dissolves  and  is  converted  into  bismuth  nitrate. 
Addition  of  ammonium  hydroxide  to  the  solution  causes  the  pre- 
cipitation of  white  bismuth  hydroxide,  which  is  filtered  off  and 
dissolved  in  the  smallest  possible  amount  of  concentrated  hydro- 
chloric acid.  The  solution  of  bismuth  chloride  so  formed  is  next 
poured  into  water,  when  bismuth  oxychloride  precipitates  as  the 
result  of  the  hydrolysis  of  the  chloride. 

EXERCISES 

1.  Write  the  equations  for  the  reactions  which  take  place  between  the 
following  substances:    (a)  P  and  HNO3,    (6)  Ca3(PO4)2  and  H2SO4,    (c)  PH3 
and  O2,    (d)  Na  and  P,    (e)  Mg  and  P,    (/)  Al  and  P,    (g)  P2O3  and  H2O, 
(A)  P2O6  and  HNO3,    (i)  P2O6  and  H2SO,,    (j)  P2OS  and  C6H10O6,    (/c)  H4P2O7 
and  H2O,    (I)  POC13  and  H2O,    (ro)  P  and  KC1O3. 

2.  How  could  you  separate    from  each  other  the  following:    (a)  P2O3 
and  P2O6,    (6)  Ca3(PO4)2  and  Na2HPO4,    (c)  Sb2S3  and  Bi2S3,    (d)  Sb?S3  and 


3.  Write  graphic  formulas  for  the  following  showing  in  each  case  the-h 
or  -   valence  of  each  element:    (a)  P2O6,    (6)  PH3,    (c)  H3PO4,    (d)  PC16. 
Write  two  possible  graphic  formulas  for    (e)  H3POS  and  two  for    (/)  H3PO2, 
and  indicate  in  each  the  valence  of  phosphorus.  , 

4.  Write  the  formulas  of  the  (a)  primary,  (6)  secondary,  and  (c)  tertiary 
calcium  salts  of  orthophosphoric  acid. 

5.  Write  an  equation  for  the  reaction  by  which  magnesium-ammonium- 
phosphate  is  formed  when  magnesium  chloride,  ammonia,  and  ammonium 
chloride  are  added  to  a  solution  of  disodium  phosphate. 

6.  Write  an  equation  for  the  reaction  by  which  PH3  and  H3PO4  are  formed 
as  the  result  of  heating  H3PO3. 

7.  What  principle  is  a  guide  in  coming  to  a  conclusion  as  to  whether  or  not 
the  compound  having  the  formula  P4S3  will  burn?     Write  an  equation  for 
the  reaction  which  takes  place  when  it  burns. 


420  INORGANIC  CHEMISTRY  FOR  COLLEGES 

8.  Describe  a  simple  experiment  which  would  show  that  the  reaction 
between  antimony  chloride  and  water  is  a  reversible  one. 

9.  An  aqueous  solution  of  I2  in  the  presence  of  NaHCO3  will  oxidize  As2O3 
to  As2O5.     (a)  Write  an  equation  for   the  reaction.     (6)  How   could   the 
reaction  be  used  to  determine  12  quantitatively. 

10.  Name  an  antidote  for  each  of  the  following:    (a)  I2,    (6)  BaCl2,    (c) 
HgCl2,  (d)  AgN03. 

11.  Write  equations  for  all  the  reactions  involved  in  the  Marsh  test  for 
arsenic. 

12.  Write  equations  for  the  reactions  which  take  place  when    (a)  Sb2O3 
is  dissolved  in  HC1  and    (6)  the  solution  is  treated  with  H2S;   when    (c)  the 
sulphide  formed  is  dissolved  in  (NH4)2S3  and  (d)  the  resulting  solution  treated 
with  HC1. 

13.  Starting  with  Ca3(PO4)2,  write  equations  for  reactions  by  which  the 
following  can  be  prepared:    (a)  H3PO4,    (6)  H3PO2,    (c)  POC13,    (d)  PH4I, 
(e)  Na4P2O7. 

14.  How  could  you  distinguish  from  each  other  the  following:  (a)  Na2HPO4 
and  Na2HAsO4,     (6)  H3PO4  and    H3PO3,     (c)  AsH3  and  SbH3,     (d)  BiOCl 
and  SbOCl,  (e)  Bi(OH)3  and  Sb(OH)3? 

15.  One  type  of  fiiction  matches  is  made  by  first  dipping  small  wooden 
sticks  in  paraffin  and  then  coating  the  ends  with  a  mixture  of  potassium 
chlorate,  glue,  powdered  flint,  and  clay.     On  the  end  of  the  tip  is  placed  a 
mixture  of  red  phosphorus  (or  P4S3),  potassium  chlorate,  glue,  and  clay.    The 
heads  of  safety  matches  may  contain  antimony  trisulphide,  in  addition  to  the 
above;   the  red  phosphorus  and  powdered  flint  are  placed  on  the  box.  Explain 
why  each  material  is  used. 


CHAPTER  XXIX 
SOME  IMPORTANT  ORGANIC  COMPOUNDS 

496.  The  compounds  of  carbon  which  were  described  in  Chap- 
ter XV  are  all  simple  in  composition  and  can  readily  be  obtained 
from  mineral  sources.  They  exhibit  properties,  in  the  main, 
like  those  of  the  analogous  substances  derived  from  the  other 
elements.  A  large  number  of  compounds  of  carbon  are  pro- 
duced as  the  result  of  the  life  processes  in  plants  and  animals  and 
from  these  many  others  have  been  made.  Carbon  compounds  of 
this  kind  are  considered  in  detail  in  a  separate  branch  of  chemistry, 
called  organic  chemistry,  because  the  principles  underlying  their 
chemical  behavior  are  quite  different  from  those  arrived  at  from 
the  study  of  inorganic  substances. 

The  complexity  in  composition  of  the  compounds  of  carbon 
results  from  the  fact  that  the  atoms  of  the  element  can  unite  with 
each  other  and  thus  form  molecules,  which  may  contain  as  many  as 
60  atoms  of  carbon.  The  formulas  of  some  of  the  simpler  hydro- 
carbons, which  are  compounds  of  hydrogen  and  carbon,  will  illus- 
trate this  fact.  The  graphic  formulas  of  methane,  CH4,  ethane, 
C2He,  and  propane,  CsHg,  are,  respectively,  as  follows: 

H  H     H  H     H    H 

H— C— H,         H— C— C— H,     and    H— C— C— C— H 

i         U  ii  A 

The  study  of  organic  compounds  has  centered  largely  around 
the  determination  of  the  way  in  which  the  atoms  are  joined  together 
in  the  molecules  of  these  substances.  When  a  number  of  atoms  are 
present  it  is  possible  for  two  substances  to  have  the  same  compo- 
sition but  entirely  different  physical  and  chemical  properties.  For 
example,  there  are  two  compounds  of  the  formula  CoHeO;  it  has 
been  shown  that  in  one  the  atoms  are  arranged  as  represented  by 

421 


422  INORGANIC  CHEMISTRY  FOR  COLLEGES 

formula    1    below,    and   in   the   other   the    arrangement   is   that 
indicated  in  formula  2: 


H    H  H  H 

I  I  I 

(1)  H— C— C— O— H  (2)  H— ( 

H 


1  i  i      i 


The  first  is  the  graphic  formula  of  grain  alcohol,  a  liquid  which 
boils  at  76°,  and  the  second  that  of  methyl  ether,  which  is  a  gas. 
Such  compounds  are  said  to  be  isomeric,  because  they  contain 
the  same  number  of  the  several  atoms.  Only  a  few  of  the  more 
important  compounds  will  be  mentioned  below. 

497.  Natural  Gas  and  Petroleum. — The  gas  that  exists  in  the 
earth  and  is  obtained  by  boring  wells  consists  chiefly  of  methane, 
CH4,  with  which  small  amounts  of  hydrogen  and  ethane,  C2He, 
are  mixed.     It  is  usually  found  in  the  localities  which  furnish 
petroleum. 

Petroleum  occurs  in  large  quantities  in  the  United  States  in 
Pennsylvania,  Ohio,  •  California,  and  Texas.  It  is  obtained  also 
in  Mexico,  certain  parts  of  Central  Europe,  Japan,  and  India. 
Petroleum  is  a  thick,  greenish  oil  which  consists  essentially  of  a 
mixture  of  a  large  number  of  hydrocarbons.  The  crude  material 
is  distilled  and  in  this  way  separated  into  products  which  are  sold 
under  trade  names  for  specific  purposes.  Gasoline,  which  is  used 
as  a  motor  fuel,  is  variable  in  composition,  but  consists  largely 
of  hydrocarbons  which  boil  between  70°  and  120°.  That  obtained 
from  most  American  petroleums  contains  the  hydrocarbons  hexane, 
CeHi4,  heptane,  CjHie,  and  octane,  CsHig.  Kerosene  consists 
of  hydrocarbons  which  boil  at  higher  temperatures  and  contain 
from  10  to  16  carbon  atoms.  Vaseline  contains  the  hydrocarbons 
of  the  formula  C22H46  and  C23H48.  Paraffin  is  a  mixture  of 
hydrocarbons  of  high  molecular  weight,  which  are  crystalline  at 
room  temperature.  The  hydrocarbons  mentioned  above  belong 
to  what  is  called  the  paraffin  series;  they  are  characterized  by 
great  inertness  to  chemical  reagents. 

498.  Unsaturated  Hydrocarbons. — A  large  number  of  hydro- 
carbons are  known  which  contain  a  smaller  proportion  of  hydrogen 
than  the  corresponding  compounds  in  the  paraffin  series;    thus, 
the  formula  of  ethane  is  CoHe,  and  that  of  ethylene  CoH4.     The 


SOME  IMPORTANT  ORGANIC  COMPOUNDS  423 

compounds  related  to  ethylene  react  readily  with  chlorine,  bro- 
mine, sulphuric  acid,  and  other  reagents.  In  order  to  indicate 
this  fact,  the  graphic  formula  of  ethylene  is  written  as  given  below 
in  the  following  equation,  which  represents  the  union  of  the 
hydrocarbon  with  bromine : 

H     H  H     H 

II  II 

C=  C  +  Br2  =  Br— C— C— Br 

II  II 

H    H  H     H 

Ethylene  is  said  to  be  unsaturated  because  it  can  unite  with 
other  substances  by  direct  addition.  This  behavior  results  from 
the  fact  that  in  the  hydrocarbon  two  carbon  atoms  are  in  combina- 
tion with  four  univalent  atoms,  whereas  the  two  carbon  atoms 
can  be  united  and  hold  in  combination  six  univalent  atoms. 

Unsaturated  hydrocarbons  are  present  in  coal  gas,  and  as  they 
burn  with  a  luminous  flame  they  are  the  so-called  illuminants. 
Acetylene,  (217),  is  still  more  highly  unsaturated;  it  has  the 
formula  HC^CH. 

499.  Carbohydrates. — Plants  are  composed  largely  of  com- 
pounds of  carbon,  hydrogen,  and  oxygen,  that  are  called  carbo- 
hydrates. Cellulose,  (CeHioOs)*,  is  the  chief  constituent  of  the 
woody  fiber  of  trees  and  other  plants.  It  exists  in  the  pure  con- 
dition in  cotton  and  linen.  Paper  is  made  from  a  pulp  prepared 
by  beating  cotton  or  linen  rags  with  water.  The  mixture  is  placed 
on  a  perforated  screen  and,  after  the  water  has  drained  off,  the 
thin  layer  of  moist  pulp  is  compressed  and  dried  by  being  passed 
through  heated  rollers.  Paper  is  also  made  in  a  similar  way 
from  wood,  after  the  latter  in  the  form  of  small  chips  has  been 
heated  under  pressure  with  a  solution  of  calcium  bisulphite  to 
dissolve  the  gums  and  substances  other  than  cellulose  in  the  wood. 
In  order  to  prevent  the  spreading  of  ink  on  paper  the  surface  is 
sized  by  covering  it  with  a  thin  layer  of  gelatine  or  rosin. 

Starch,  (CeHioOs)!,,  is  obtained  from  corn,  wheat,  potatoes, 
and  other  plants.  It  consists  of  minute  granules  which  have  a 
characteristic  appearance  when  examined  with  a  microscope. 
The  granules  are  covered  with  a  very  thin  coating  of  a  form  of 
cellulose.  When  starch  is  heated  with  water  the  coatings  burst, 
and  when  the  mixture  is  cooled  a  jelly  is  obtained.  A  part  of  the 


424  INORGANIC  CHEMISTRY  FOR  COLLEGES 

starch  passes  into  colloidal  solution  (534).  When  a  dilute  solu- 
tion of  iodine  comes  into  contact  with  starch  that  has  been  heated 
with  water,  a  characteristic  blue  color  is  formed,  which  serves  as  a 
test  for  either  iodine  or  starch  (412).  Wheat  flour  consists  of 
starch  and  a  small  proportion  of  gluten,  which  is  a  mixture  of 
certain  nitrogenous  compounds  that  are  known  as  proteins. 
Gluten  is  essential  in  bread-making,  because  it  is  sticky  and,  there- 
fore, prevents  the  escape  of  the  gas  produced  when  dough  rises. 

Glucose,  CeH^Oe,  is  present  in  grapes,  and  for  this  reason  is 
sometimes  called  grape  sugar.  It  is  also  present  in  honey,  and  is 
formed  when  starch  is  heated  with  water;  if  a  trace  of  an  acid  is 
present  the  change  takes  place  much  more  quickly.  Glucose  is 
made  in  this  way  from  starch,  and  after  the  neutralization  of  the 
acid  is  sold  in  the  form  of  a  syrup. 

Sucrose,  Ci2H22On,  is  obtained  from  the  sugar  cane  or  sugar 
beets;  it  is  commonly  called  cane  sugar.  When  sugar  is  heated 
with  water  in  the  presence  of  a  small  amount  of  an  acid  or  any 
substance  which  furnishes  hydrogen  ions,  it  hydrolyzes  and  glu- 
cose and  fructose,  another  sugar  present  in  honey  and  certain  fruits, 
are  formed : 

C12H22On  +  H20  =  C6Hi206  +  C6H1206 

The  two  simpler  sugars  contain  the  same  number  of  carbon, 
hydrogen,  and  oxygen  atoms;  they  differ  in  the  structure  of  their 
molecules,  that  is,  in  the  way  in  which  the  atoms  are  joined  to  one 
another.  This  change  of  sugar  into  glucose  and  fructose  is  brought 
about  in  candy-making  by  heating  sugar  with  cream  of  tartar 
(501),  which  is  an  acid  salt.  The  mixture  formed  does  not  crys- 
tallize readily  and,  therefore,  gives  the  candy  a  smooth  texture. 

Lactose,  Ci2H22On,  is  a  sugar  which  is  obtained  from  milk; 
the  latter  contains,  in  addition  to  lactose,  fat  and  casein,  which 
is  a  protein. 

500.  Alcohols. — When  yeast  is  put  into  aqueous  solutions  of 
most  sugars  fermentation  takes  place,  and  alcohol  and  carbon 
dioxide  are  formed.  The  reaction  in  the  case  of  glucose  is  as  fol- 
lows: 

C6Hi2O6  =  2C2H6O  +  2CO2 

Yeast  contains  substances  which  change  the  more  complex  car- 
bohydrates into  glucose  and,  as  a  consequence,  fermentation  takes 


SOME  IMPORTANT  ORGANIC  COMPOUNDS  425 

place  when  yeast  is  mixed  with  moist  flour.  It  is  in  this  way  that 
the  dough  for  bread-making  is  prepared. 

Alcohol  is  obtained  by  fermentation  of  the  carbohydrates  in 
grains.  On  account  of  the  fact  that  it  reacts  with  strong  acids  to 
form  compounds  as  the  result  of  the  replacement  of  a  hydrogen 
and  an  oxygen  atom,  its  formula  is  usually  written  C2HsOH. 
The  radical  C2H.5  occurs  in  a  large  number  of  compounds  and  has 
been  given  the  name  ethyl;  for  this  reason  grain  alcohol  is  called 
ethyl  alcohol.  Alcohol  boils  at  76°  and  burns  with  a  light  blue 
flame. 

Methyl  alcohol,  CHaOH,  is  formed  when  wood  is  heated  to 
make  charcoal;  for  this  reason  it  is  often  called  wood  alcohol.  It 
is  obtained  by  condensing  the  vapors  given  off  and  distilling  the 
resulting  liquid.  Methyl  alcohol  is  very  poisonous  and  produces 
blindness.  When  it  burns,  carbon  dioxide  and  water  are  formed, 
but  if  an  insufficient  amount  of  air  is  used  partial  oxidation  takes 
place  and  formaldehyde,  CEbO,  is  produced.  The  latter  is  a  gas 
which  dissolves  in  water.  A  40  per  cent  solution  of  formaldehyde 
in  water  is  sold  under  the  name  formaline,  and  is  used  as  a  disin- 
fectant. 

501.  Organic  Acids. — The  acids  derived  from  carbon  have  the 

O 

grouping  of  elements  represented  by  the  formula  — C — OH  (or 
COOH),  which  is  called  the  carboxyl  group.  The  formula  of 

O 

carbonic  acid  represented  in  this  way  is  HO — C — OH.  Acetic 
acid  has  the  formula  CHs-COOH.  It  is  obtained  along  with 
methyl  alcohol  when  wood  is  distilled.  The  acid  is  also  formed 
as  the  result  of  fermentation  when  fruit  juices  are  exposed  to  the 
air,  for  the  latter  under  ordinary  circumstances  contains  micro- 
organisms which  convert  the  sugars  present  into  acetic  acid.  It  is 
in  this  way  that  vinegar  is  formed  from  the  juice  of  apples  or  from 
wine.  The  organisms  which  bring  about  the  change  are  present 
in  what  is  called  mother  of  vinegar.  Lactic  acid,  C2H5O-COOH, 
is  formed  when  the  sugar  in  milk  ferments;  it  is  present  in  sour 
milk. 

The  salts  of  a  large  number  of  acids  are  present  in  fruits. 
Tartaric  acid  is  a  dibasic  acid  obtained  from  grapes;  it  has  the 


426  INORGANIC  CHEMISTRY  FOR  COLLEGES 


formula  C2H402(COOH)2  (or  H2  -CJ^Oe);  cream  of  tartar  is 
the  acid  potassium  salt.  Citric  acid,  CsHsCXCOOH^,  is  obtained 
from  lemons,  benzoic  acid,  CeHs-COOH,  from  cranberries,  and 
oxalic  acid,  (COOH)2  (or  H2C2O4),  from  rhubarb  and  other  plants. 
Oxalic  acid  forms  a  dihydrate,  H2C2C>4,2H2O,  when  crystallized 
from  water.  It  is  manufactured  by  heating  sawdust  with  sodium 
hydroxide  and  a  small  amount  of  water.  From  the  sodium  salt 
formed  in  this  way,  the  free  acid  is  obtained  by  treatment  with 
sulphuric  acid. 

502.  Esters.  —  The  products  formed  as  the  result  of  the  inter- 
action of  alcohols  and  acids  are  called  esters;  ethyl  acetate,  for 
example,  is  the  ester  formed  from  ethyl  alcohol  and  acetic  acid; 
it  has  the  formula  CH3-COOC2H5. 

CHs-COOH  +  C2H5OH  —  CH3-COOC2H5  +  H2O 

The  reaction  indicated  is  reversible,  for  when  an  ester  is  heated 
with  water  it  is  converted  into  an  acid  and  an  alcohol.  The 
pleasant  odor  of  many  plants  results  from  the  presence  of  esters. 

Fats  are  mixtures  of  esters  of  acids  of  high  molecular  weight. 
The  alcohol  obtained  from  them  when  they  are  heated  with 
water  or  an  alkali  is  glycerine,  CsH^OH^.  This  reaction  is 
what  occurs  in  soap  making.  Among  the  esters  in  beef  fat 
are  the  glycerine  esters  of  palmitic  acid,  C 
stearic  acid,  CiyHss-COOH,  and  oleic  acid,  C 
which  is  an  unsaturated  acid.  When  the  ester  of  stearic  acid,  which 
has  the  formula  (CiyH^-COO^CsHs,  is  heated  with  a  solution 
of  sodium  hydroxide,  sodium  stearate,  CiyHss-COONa,  and 
glycerine  are  formed.  The  mixture  of  salts  produced  in  this  way 
from  fats  is  soap.  The  behavior  of  soap  with  water  and  its  use 
as  a  cleansing  agent  are  described  later  (628,  629). 

Fats  from  different  sources  differ  from  one  another  in  the  esters 
present  and  their  relative  amounts.  Butter  is  characterized  by 
containing  about  5  per  cent  of  the  ester  of  butyric  acid, 
CsHT-COOH.  The  liquid  fats,  such  as  olive  oil,  contain  a  high 
percentage  of  esters  derived  from  an  unsaturated  acid.  Stearin, 
(CiTHss'COO^CsHs,  which  occurs  in  beef  tallow,  is  a  solid,  and 
olein,  (Ci7H33-COO)3C3H5,  which  is  present  in  edible  oils,  is  a 
liquid  at  ordinary  temperatures. 


SOME  IMPORTANT  ORGANIC  COMPOUNDS  427 

Cellulose,  which  is  an  alcohol,  forms  an  ester  when  treated 
with  nitric  acid;  cellulose  trinitrate  is  used  as  an  explosive  (364). 
A  mixture  of  camphor  and  the  nitrates  of  cellulose  which  contain 
less  nitrogen  than  the  trinitrate  is  called  celluloid.  The  nitrate 
made  from  glycerine,  CsH^NOa^,  is  the  important  constituent 
of  dynamite  (364). 


CHAPTER  XXX 

SILICON  AND   BORON.    THE  ACID-FORMING  ELEMENTS  AND 
THE  PERIODIC  CLASSIFICATION 

503.  Silicon  is  the  second  member  of  the  carbon  family,  which 
is  in  the  fourth  group  in  the  periodic  classification.     It  resembles 
carbon  closely  in  many  chemical  properties,  but  is  more  active 
toward  electronegative  elements;   it  unites  readily  with  chlorine, 
for  example,  whereas  carbon  does  not  react  directly  with  the 
halogen.     Its  greater  activity  is  also  shown  by  the  fact  that  when 
sodium  carbonate  is  heated  to  a  high  temperature  with  silicon, 
sodium  silicate  is  formed  and  carbon  is  set  free. 

Silicon  has  the  valence  4  in  its  compounds  with  either  positive 
or  negative  elements;  the  only  compound  in  which  it  does  not 
show  this  valence  is  silicon  monoxide,  SiO;  it  always  acts  as  an 
acid-forming  element.  The  consideration  of  tin  and  lead,  two 
important  elements  in  the  fourth  group  of  the  periodic  classifica- 
tion, will  be  deferred  until  later  on  account  of  the  fact  that  their 
most  characteristic  properties  are  those  of  metals. 

504.  Occurrence. — Next  to  oxygen,  silicon  is  the  most  abundant 
element  in  the  earth's  crust,  of  which  it  constitutes  about  26  per 
cent.     It  is  never  found  free,  but  in  combination  with  oxygen  as 
silicon  dioxide,  an  impure  variety  of  which  is  sand,  or  in  the  form 
of  silicates,  which  are  salts  derived  from  the  acid  EbSiOa,  and  are 
the  chief  constituents  of  rocks  and  clays.     The  soil  is  a  mixture 
of  sand,  silicates,  and  the  complex  organic  compounds  formed  as 
the  result  of  the  decomposition  of  dead  vegetable  material.     The 
silicates  that  occur  in  nature  and  those  made  artificially  have 
extensive  uses  on  account  of  their  physical  properties. 

505.  Preparation. — Silicon  is  obtained  by  heating  a  mixture  of 
sand  and  carbon  in  an  electric  furnace  (218);    carbon  monoxide 
and  silicon  are  the  chief  products  of  the  reaction,  but  some  silicon 
monoxide  is  produced  in  the  parts  of  the  furnace  which  are  at  a 

428 


SILICON  AND  BORON  429 

lower  temperature  than  the  central  core  through  which  the  major 
part  of  the  current  passes.  Ferrosilicon,  an  alloy  of  iron  and  sili- 
con, is  prepared  on  the  large  scale  by  a  similar  process  from  a  mix- 
ture of  oxide  of  iron,  sand,  and  carbon;  it  is  used  in  the  manu- 
facture of  steel.  Silicon  is  an  important  constituent  of  the  crude 
iron  obtained  from  blast-furnaces  (pig-iron),  and  has  a  marked 
effect  on  its  properties. 

In  the  laboratory,  silicon  is  prepared  by  igniting  a  mixture  of 
powdered  silicon  dioxide  and  magnesium.  It  is  necessary  only  to 
start  the  reaction,  which  is  a  vigorous  one: 

SiO2  +  2Mg  =  2MgO  +  Si 

The  silicon  is  obtained  by  treating  the  mixture  with  an  acid, 
which  dissolves  the  magnesium  oxide  and  the  small  amount  of 
magnesium  silicide,  Mg2Si,  present.  Silicon  can  be  obtained  in  a 
crystalline  condition  by  heating  it  with  molten  zinc,  and  after 
the  mass  has  solidified,  dissolving  out  the  metal  with  an  acid. 

506.  Properties. — Amorphous  silicon,  which  is  a  brown  powder, 
is  more  active  than  the  crystalline  variety;  at  ordinary  tempera- 
tures it  unites  with  fluorine  and,  when  heated,  with  chlorine,  bro- 
mine, oxygen,  and  sulphur.     At  high  temperatures  it  combines 
with  nitrogen,  and  at  the  temperature  of  the  electric  furnace  with 
carbon  and  boron.     It  is  slowly  oxidized  by  aqua  regia  to  silicic 
acid,  and  is  dissolved  by  a  mixture  of  hydrofluoric  and  nitric  acids. 
Silicon  forms  silicides  with  the  more  active  metals. 

Silicon  reacts  with  solutions  of  the  alkalies: 

2NaOH  +  Si  +  H2O  =  Na2SiO3  +  2H2 

The  reaction  is  utilized  in  making  hydrogen  for  filling  balloons. 

Crystalline  silicon  is  obtained  in  the  form  of  black  needles, 
which  are  harder  than  glass  and  will,  therefore,  scratch  it.  In 
order  to  bring  it  into  reaction  a  higher  temperature  is  required  than 
is  the  case  with  the  amorphous  element. 

507.  Silicon  Dioxide. — Silicon    dioxide,  which  is  also   called 
silica,  exists  in  nature  widely  distributed  in  a  more  or  less  pure 
condition  as  sand,  which  is  ordinarily  colored  as  the  result  of  the 
presence  of  compounds  containing  iron.     Sandstone  is  made  up 
of  sand  grains  cemented  together  by  amorphous  silica,  clay,  and 
other  substances;  it  is  extensively  used  for  building  purposes. 


430  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Quartz  or  rock  crystal  is  the  name  given  to  a  crystalline  variety 
of  silicon  dioxide,  which  occurs  as  clear,  colorless,  six-sided  prisms 
terminated  by  pyramids.  Many  varieties  of  quartz  are  found 
that  are  colored  as  the  result  of  the  presence  of  small  amounts  of 
impurities;  among  these  are  amethyst,  which  contains  traces  of 
manganese  and  iron,  and  yellow,  rose,  and  smoky  quartz,  and  cat's 
eye.  If  the  silica  is  not  crystalline  and  is  colored  brown  or  red 
by  iron  oxide,  it  is  known  as  agate,  onyx,  or  jasper.  Opal  and  flint 
are  varieties  of  hydrated  silica.  Infusorial  or  diatomaceous  earth, 
which  is  also  called  Tripoli  powder,  is  a  form  of  silica  made  up  of 
the  skeletons  of  minute  organisms.  On  account  of  its  physical 
condition  it  is  used  as  a  polishing  powder,  and  on  account  of  its 
absorbent  qualities  for  removing  the  coloring  matter  from  oils. 

Rock  crystal  is  used  in  making  special  kinds  of  chemical  appa- 
ratus required  when  reactions  are  to  be  carried  out  at  very  high 
temperatures  or  when  a  vessel  must  be  used  that  is  not  attacked 
by  acids.  Since  quartz  melts  at  1600°  the  apparatus  must  be 
made  with  the  aid  of  the  oxy-hydrogen  flame  or  in  an  electric  fur- 
nace. Apparatus  made  from  rock  crystal  is  transparent  and 
resembles  glass  in  appearance;  it  possesses  one  property  that  is 
very  valuable — it  expands  to  such  a  slight  degree  when  heated 
that  it  can  undergo  great  changes  in  temperature  without  crack- 
ing; a  quartz  vessel  can  be  heated  red-hot  and  plunged  into  cold 
water  without  breaking.  Apparatus  made  from  other  varieties 
of  silicon  dioxide  are  milky  in  appearance,  but  on  account  of  their 
lower  cost  are  used  in  place  of  that  prepared  from  rock  crystal 
when  transparency  is  not  required.  It  has  recently  been  found 
possible  to  construct  silica  apparatus  of  such  a  size  that  it  can  be 
used  in  chemical  operations  carried  out  on  the  industrial  scale. 
Quartz  is  much  more  transparent  to  ultra-violet  light  than  glass 
and  is  used  in  place  of  the  latter  in  experimental  work  with  this 
form  of  energy. 

508.  Silicic  Acid  and  Silicates. — Although  silicon  dioxide  does 
not  react  with  water  to  form  an  acid,  it  possesses  the  most  char- 
acteristic property  of  an  acid  anhydride — it  dissolves  in  solutions 
of  the  alkalies  to  form  salts,  which  are  known  as  silicates.  As 
the  reaction  takes  place  very  slowly,  sodium  silicate,  which  has 
extensive  uses,  is  made  by  fusing  sand  with  sodium  hydroxide  or 
sodium  carbonate.  When  an  acid  is  added  to  a  solution  of  sodium 


SILICON  AND  BORON  431 

silicate,  silicic  acid  is  formed  and  appears  as  a  gelatinous  precipitate. 
The  composition  of  the  precipitate  is  approximately  that  repre- 
sented by  the  formula  of  orthosilicic  acid,  H4SiO4.  When  it  is 
dried  it  slowly  loses  water  until  silicon  dioxide  is  formed.  A  large 
number  of  salts  are  known  derived  from  metasilicic  acid,  H2SiO3, 
which  resembles  carbonic  acid  in  composition. 

509.  Many  minerals  are  derived  from  what  are  called  poly- 
silicic  acids,  which  are  formed  as  the  result  of  the  partial  dehy- 
dration of  two  or  more  molecules  of  orthosilicic  acid;  they  may  be 
considered  as  made  up  of  silicon  dioxide  in  combination  with  water 
in  different  proportions,   the  relation  between  the  acids  being 
similar  to  that  already  seen  in  the  case  of  the  phosphoric  acids. 

If  1  molecule  of  orthosilicic  acid  loses  1  molecule  of  water,  it  is 
transformed  into  metasilicic  acid, 

H4SiO4  -  H2O  =  H2SiO3  (H2O,SiO2) 

Polysilicic  acids  are  derived  from  two  or  more  molecules  of  the  ortho 
acid;  the  more  important  of  these  contain  2  and  3  atoms  of  silicon. 
The  salts  of  two  disilicic  acids  are  known ;  one  acid  is  formed  from 
2  molecules  of  the  ortho  acid  by  the  loss  of  1  molecule  of  water, 

2H4SiO4  -  H2O  =  H6Si2O7  (3H2O,2SiO2) 
and  the  other  by  the  loss  of  3  molecules, 

2H4SiO4  -  3H2O  =  H2Si2O5  (H2O,2SiO2) 

Trisilicates  are  known  which  are  derived  from  an  acid  formed  as 
the  result  of  the  loss  of  4  molecules  of  water  from  3  molecules  of 
orthosilicic  acid, 

3H4SiO4  -  4H2O  =  H4Si3O8  (2H2O,3SiO2) 

510.  A  large  number  of  minerals  are  salts  of  these  four  silicic 
acids,  but  only  a  few  will  be  mentioned.    Garnet,  which  is  some- 
times used  as  a  jewel,  is  an  orthosilicate  of  magnesium  and  alumin- 
ium in  which  more  or  less  of  the  magnesium  is  replaced  by  calcium. 
To  express  this  latter  fact  the  symbols  (Mg,Ca)  are  written  in 
the  formula,  which  is  (Mg,Ca)3Al2(SiO4)3.     Mica  is  a  familiar 
mineral  that  is  used  for  a  number  of  purposes  on  account  of  its 
peculiar  physical  structure  which  allows  it  to  be  separated  into 
thin  sheets;    it   is   also  an  orthosilicate  and  has   the  formula 


432  INORGANIC  CHEMISTRY  FOR  COLLEGES 

(K,Na)H2Al3  (8104)3:  it  is  extensively  used  in  the  construction  of 
electrical  apparatus  because  it  does  not  conduct  electricity  and 
resists  fairly  high  temperatures.  Kaolin,  clay,  is  a  hydrated 
aluminium  orthosilicate,  H2Al2(SiO4)2,H2O.  Serpentine  is  an 
example  of  a  disilicate  and  has  the  formula  Mg3Si2O7,2H2O. 
Feldspar,  or  orthoclase,  which  plays  an  important  part  in  soil- 
formation,  and  is  used  in  making  porcelain  and  other  table- 
ware, is  a  trisilicate  and  has  the  formula  KAlSisOg. 

A  large  number  of  metasilicates  are  important  minerals. 
Among  these  are  talc  (soapstone),  H2Mg3 (810)3)4,  and  asbestos, 
MgsCa  (8163)4.  The  minerals  that  occur  as  silicates  are  the  chief 
constituents  of  many  igneous  rocks,  which  are  usually  mechanical 
mixtures  of  two  or  more  minerals.  Granite,  for  example,  which  is  a 
very  widely  distributed  rock,  contains  feldspar,  mica,  and 
quartz. 

511.  Sodium  silicate  is  manufactured  on  a  large  scale,  as  it 
has  important  industrial  applications;  it  is  known  as  water-glass 
on  account  of  the  fact  that  it  is  soluble  in  water,  and  is  obtained 
as  a  transparent  amorphous  mass  resembling  glass  when  its  solu- 
tion is  evaporated  to  dry  ness.  It  is  made  by  fusing  together  for 
eight  to  ten  hours  at  a  high  temperature  a  mixture  of  quartz  or 
infusorial  earth  and  caustic  soda,  sodium  carbonate,  or  sodium  sul- 
phate. The  product  is  a  glass  which  is  powdered  and  heated 
with  water  under  pressure  until  the  soluble  material  is  dissolved. 
The  solution  is  then  evaporated  until  it  has  the  specific  gravity  1.7. 
The  composition  of  the  silicate  formed  is  determined  by  the 
proportions  of  the  ingredients  used;  it  approximates  that  of  the 
formula  Na2Si4Og  (Na2Si03,3SiO2).  It  is  used  as  a  fixative  for 
pigments  in  calico  printing,  for  rendering  cloth  and  paper  non- 
inflammable,  for  preserving  eggs,  in  cement  mixtures  for  glass, 
wood,  and  leather,  as  a  preservative  for  timber  and  porous  stone, 
in  the  manufacture  of  artificial  stone,  and  as  a  size  for  paper  and 
fabrics.  Sodium  silicate  is  mixed  with  soap  in  the  preparation  of 
certain  varieties  of  the  latter  used  for  laundry  work,  because  it 
serves  the  valuable  purpose  of  softening  hard-water. 

All  the  silicates  are  insoluble  except  those  of  the  alkali  metals. 
Since  silicic  acid  is  a  very  weak  acid,  its  soluble  salts  are  hydro- 
lyzed  by  water  and  show  an  alkaline  reaction.  The  use  of  sodium 
silicate  as  a  paint-remover  is  based  on  this  fact;  the  alkali  formed 


SILICON  AND  BORON  433 

as  the  result  of  the  hydrolysis  of  the  salt  attacks  the  solidified  oil 
in  the  paint  and  converts  it  into  substances  soluble  in  water. 

512.  Test  for  Silicates. — When  an  acid  is  added  to  a  solution  of 
a  silicate,  silicic  acid  is  precipitated  in  a  gelatinous  condition.     To 
complete  the  identification  of  the  latter  it  is  filtered  off,  treated 
with  hydrofluoric  acid,  and  evaporated  to  dryness;    silicic  acid 
leaves  no  residue,  as  it  is  converted  by  hydrofluoric  acid  into  silicon 
fluoride,  which  is  volatile.     If  the  substance  thought  to  be  a  sili- 
cate is  insoluble,  it  is  fused  at  red-heat  with  sodium  carbonate, 
which  converts  insoluble  silicates  into  soluble  sodium  silicate, 
and  this  is  tested  as  just  described. 

513.  Hydrides  of  Silicon. — Silicon  forms  hydrides  of  the  com- 
position SiH4,  Si2He,  and  Si2H2.     They  are  all  gases  which  burn 
in  air.     The  tetrahydride  can  be  prepared  by  treating  magnesium 
silicide  with  hydrochloric  acid : 

Mg2Si  +  4HC1  =  2MgCl2  +  SiH4 

When  prepared  in  this  way  it  contains  an  impurity  which  renders 
it  spontaneously  inflammable;  it  is  readily  decomposed  by  heat 
into  its  constituents. 

514.  Halides   of   Silicon. — Silicon   tetrachloride   can   be   pre- 
pared by  the  direct  union  of  the  elements,  or  if  silicon  is  not  avail- 
able, by  the  action  of  chlorine  and  carbon  at  a  high  temperature 
on  silicon  dioxide: 

SiO2  +  2C  +  2C12  =  SiCl4  +  2CO 

Neither  carbon  nor  chlorine  alone  will  react  with  silica  at  the 
temperature  which  can  be  used  to  prepare  the  chloride  in  this  way. 
In  the  simultaneous  action  of  carbon  and  chlorine  the  affinity  of 
the  former  for  oxygen  and  the  latter  for  silicon  come  into  play  at 
the  same  time  and  the  change  takes  place. 

Silicon  tetrachloride  is  a  colorless  liquid,  which  boils  at  59° 
and  fumes  in  the  air  as  the  result  of  the  hydrolysis  brought  about 
by  the  moisture  present: 

SiCl4  +  4H2O  =  Si(OH)4  +  4HC1 

When  the  vapor  of  the  chloride  comes  in  contact  with  air  contain- 
ing ammonia  a  very  dense  cloud  is  formed.  The  reaction  was 


434  INORGANIC  CHEMISTRY  FOR  COLLEGES 

utilized  in  the  recent  war  in  producing  smoke  clouds  to  hide 
vessels  at  sea. 

Silicon  tetrafluoride  (426)  can  be  prepared  by  heating  together 
a  mixture  of  calcium  fluoride,  sand,  and  concentrated  sulphuric 
acid;  hydrofluoric  acid  is  first  formed,  and  then  reacts  with  the 
silicon  dioxide  to  produce  silicon  fluoride  and  water.  Silicon 
tetrafluoride  is  a  gas,  which  reacts  vigorously  with  water: 

SiF4  +  2H2O  ^±  Si(OH)4  +  2H2F2 

The  hydrofluoric  acid  produced  unites  with  some  of  the  silicon 
fluoride  present  and  forms  hydrofluosilicic  acid, 

SiF4  +  H2F2  =  H2SiF6 

which  can  be  considered  as  derived  from  silicic  acid,  H2Si03,  as 
the  result  of  the  replacement  of  3  oxygen  atoms  by  6  fluorine  atoms. 
Hydrofluosilicic  acid  is  stable  in  solution  only.  When  a  solu- 
tion is  evaporated,  silicon  fluoride  is  given  off.  The  acid  can  be 
used  for  testing  for  potassium,  as  its  salt  containing  this  metal  is 
difficultly  soluble  in  water. 

BORON 

515.  Boron  is  the  first  member  in  the  third  group  in  the  periodic 
classification  of  the  elements.     It  exhibits  in  most  of  its  compounds 
the  properties  of  an  acid-forming  element,  but  it  forms  an  acid  sul- 
phate and  a  phosphate,  which  are  readily  hydrolyzed  by  water. 
The  second  member  of  the  group,  aluminium,  is  a  well-charac- 
terized metal. 

The  compounds  of  boron  which  occur  in  nature  and  those 
which  have  important  uses  are  derivatives  of  boric  acid.  The 
element  has  a  strong  affinity  for  oxygen  and  in  the  course  of  the 
formation  of  the  earth  united  with  this  element;  since  the  oxide 
is  an  acid  anhydride,  salts  were  produced  as  the  result  of  combina- 
tion taking  place  between  it  and  metallic  oxides. 

516.  Occurrence. — Boric  acid,  HaBOs,  is  a  very  weak  acid,  and 
like  other  weak  oxygen  acids  it  forms  salts  derived  from  acids  the 
molecules  of  which  contain  a  number  of  the  atoms  of  the  acid- 
forming  element.     The  polyboric  acids  bear  a  relation  to  normal 
boric  acid,  HsBOs,  similar  to  that  which  exists  between  silicic  acid, 
H4Si04,  and  the  polysilicic  acids. 


SILICON  AND  BORON  435 

Three  important  minerals  containing  boron  are  borax, 
(Na2O,2B2O3),  which  is  a  salt  of  tetraboric  acid;  cole- 
manite,  Ca2B6Oii,5H2O  (2CaO,3B2O3,5H2O) ;  and  boracite 
Mg3B8Oi5,MgCl2  (3MgO,4B2O3,MgCl2).  Since  the  salts  of 
boric  acid  are  hydrolyzed  by  water,  the  free  acid  is  formed  in 
nature  in  certain  hot  springs;  those  of  Tuscany  are  used  as  a 
commercial  source  of  the  acid. 

517.  Preparation. — The  method  used  to  isolate  boron  is  analo- 
gous to  that  employed  to  obtain  free  silicon;   its  oxide  is  heated 
with  magnesium,  and  the  resulting  mass  treated  with  an  acid  to 
dissolve  the  magnesium  oxide  formed.     The  product  is  amor- 
phous boron,  which  can  be  obtained  as  crystals  by  dissolving  it  in 
molten  aluminium  and  removing  the  latter,   after  cooling,  by  means 
of  hydrochloric  acid.     Aluminium  can  also  be  used  to  reduce  boron 
oxide,  and  if  an  excess  of  the  metal  is  used  the  product  obtained 
directly  is  crystalline  boron. 

518.  Properties. — Amorphous  boron  is  a  black  powder,  which 
can  be  fused  at  the  temperature  obtained  in  an  electric  furnace. 
It  is  brittle,  very  hard,  and  a  poor  conductor  of  heat  and  elec- 
tricity at  ordinary  temperatures.     With  rise  in  temperature  its 
physical  properties  change  rapidly;    at  400°  its  electrical  con- 
ductivity is  over  two  million  times  that  shown  at  room-tempera- 
ture, and  it  approaches  a  metal  in  electrical  properties. 

At  high  temperatures  boron  is  a  very  active  element,  which 
resembles  silicon  and  carbon  in  this  respect,  but  it  is  more  active 
than  these  elements,  for  it  will  decompose  both  silicon  dioxide 
and  carbon  monoxide.  When  boron  is  heated  with  nitrogen,  it 
forms  a  nitride,  BN,  which  is  converted  by  steam  into  boric  acid, 
H3BO3,  and  ammonia.  At  the  temperature  of  the  electric  furnace 
boron  and  carbon  form  a  carbide,  BeC,  which  is  characterized  by 
its  great  hardness,  as  it  stands  next  to  the  diamond  in  this  respect. 

Free  boron  is  added  in  small  quantities  to  copper  when  the 
latter  is  to  be  used  for  making  castings  because  it  gives  a  product 
which  is  free  from  blow-holes. 

Boron  dissolves  in  molten  potassium  hydroxide: 

2B  +  6KOH  =  2K3BO3  +  3H2 

519.  Boric  Acid. — The  acid  is  found  free  in  nature  as  has  been 
indicated.     In  Tuscany,  ponds  are  formed  around  the  fumeroles 


436  INORGANIC  CHEMISTRY  FOR  COLLEGES 

from  which  steam  issues  carrying  boric  acid  (suffioni) ;  the  steam  is 
condensed  in  the  water,  and  a  solution  obtained  which  contains 
about  2  per  cent  of  the  acid.  This  is  evaporated  by  the  steam  from 
the  fumeroles  and  crystalline  boric  acid  is  obtained. 

Boric  acid,  H3BO3,  is  made  in  California  and  in  Chile,  by 
heating  calcium  borate  (colemanite)  suspended  in  water,  through 
which  sulphur  dioxide  is  passed;  the  sulphurous  acid  formed  con- 
verts the  calcium  borate  into  boric  acid  and  calcium  bisulphite. 

Boric  acid  crystallizes  from  water  in  pearly  white,  thin  plates. 
It  is  soluble  in  25  parts  of  water  at  19°  and  in  3  parts  at  100°. 
When  heated,  it  loses  water;  at  140°  tetraboric  acid,  H^B^y,  is 
formed  and  at  red  heat  it  is  converted  into  its  anhydride,  B9Os, 
which  is  non- volatile  at  high  temperatures.  Both  tetraboric 
acid  and  boron  trioxide  react  slowly  with  water  to  form  boric  acid. 
Boric  acid  is  a  very  weak  acid;  its  aqueous  solution  barely  affects 
litmus  paper;  it  is  slowly  volatile  with  steam. 

Boric  acid  is  used  in  the  preparation  of  borax,  and  in  making 
enamels  and  glazes  for  pottery  and  in  certain  kinds  of  glass.  Its 
solution  has  antiseptic  properties  and  is  used  as  an  eye-wash  and 
for  other  purposes  in  medicine  under  the  name  boracic  acid.  It 
is  also  used  for  preserving  meat,  fish,  oysters,  and  milk,  although 
its  use  for  this  purpose  is  prohibited  by  law  in  certain  States. 

520.  Berates. — The  only  salt  derived  from  boric  acid  which 
is  used  extensively  is  borax,  sodium  tetraborate,  ^26467 .  It  is 
obtainable  commercially  as  the  pentahydrate,  ^26467, 5H2O,  or 
as  the  decahydrate,  Na2B4O7,10H2O.  Borax  is  found  native  in 
Thibet,  Ceylon,  and  California.  It  was  formerly  prepared  in 
large  quantities  from  the  water  of  Borax  Lake,  California,  by 
evaporating  it  to  crystallization.  It  is  now  obtained  from  dry 
lake  beds  in  the  Death  Valley  region,  where  the  surface  of  the 
earth  is  covered  with  a  crust  composed  of  borax  and  the  sulphate, 
chloride,  and  carbonate  of  sodium.  Most  of  the  borax  pro- 
duced in  California  is  made  from  colemanite  or  from  ulexite, 
NaCaBsOojSH^O.  The  mineral  in  either  case  is  roasted,  and  as 
the  result  of  the  loss  of  water  falls  to  a  powder;  this  is  then  boiled 
with  a  solution  of  sodium  carbonate  and  sodium  bicarbonate;  the 
calcium  is  precipitated  as  carbonate,  and  borax  is  obtained  from 
the  solution  by  evaporation. 

The  decahydrate  of  borax  is  obtained  when  the  salt  crystallizes 


SILICON  AND  BORON  437 

from  solution  at  27°  or  below;  it  forms  large  efflorescent  prisms, 
which  melt  in  the  water  of  hydra tion  when  heated,  and  swell  as  the 
latter  is  given  off,  forming  a  spongy  mass  that  at  higher  tempera- 
tures melts  to  a  clear  glass.  If  a  solution  of  borax  is  concentrated 
so  that  crystals  are  formed  at  56°  or  higher,  the  pentahydrate  is 
obtained  in  the  form  of  octahedral  crystals,  which  are  permanent 
in  dry  air;  it  fuses  without  swelling  and  is,  therefore,  preferred 
to  the  pentahydrate  when  used  as  a  flux. 

Borax  is  hydrolyzed  to  the  extent  of  about  one-half  per  cent  in 
tenth-normal  solution,  and  shows,  accordingly,  an  alkaline  reac- 
tion. Some  of  the  uses  to  which  borax  is  put  depend  upon  this 
fact.  It  is  employed,  for  example,  in  ungumming  raw  silk;  the 
small  amount  of  alkali  formed  in  the  hydrolysis  serves  to  dissolve 
the  gum  from  the  fiber,  which  would  be  seriously  affected  if  a 
solution  of  sodium  hydroxide  were  used.  The  use  of  borax  in 
soap  is  based  in  part  upon  the  fact  that  the  alkali  formed  from 
it  markedly  assists  in  converting  fat  and  oily  substances  into 
such  a  finely  divided  condition  (an  emulsion)  that  they  can  be 
readily  removed  by  water.  If  soap  is  to  be  used  with  hard  water 
the  presence  of  borax  is  advantageous  since  the  soluble  calcium 
salts  present  in  the  water  are  precipitated  as  calcium  borate  and 
do  not,  therefore,  interfere  with  the  action  of  the  soap.  Borax 
is  used  for  preserving  meat,  as  a  mordant  in  dyeing,  in  medicine 
and  pharmacy,  as  a  flux,  and  in  the  manufacture  of  enamels  and 
glazes  for  metal  ware  and  pottery. 

521.  A  flux  is  a  substance  which  converts  compounds  infusible 
at  a  certain  temperature  into  others  which  melt  at  this  tem- 
perature. Fluxes  are  generally  used  when  reactions  are  carried  out 
with  refractory  substances  at  high  temperatures,  on  account  of  the 
fact  that  the  chemical  reaction  takes  place  only  when  the  substances 
involved  are  in  intimate  contact — a  condition  difficult  to  bring 
about  in  the  case  of  solids.  The  use  of  borax  as  a  flux  is  based  upon 
the  fact  that  it  contains  an  excess  of  boric  anhydride  over  that 
required  to  form  sodium  metaborate,  NaB(>2.  The  relation  in  com- 
position between  the  two  salts  is  clearly  seen  by  comparing  their  two 
formulas  written  in  such  a  way  as  to  indicate  the  amounts  of  boron 
oxide  present;  sodium  metaborate  is  Na20,B203,  and  borax  is 
Na2O,2B2C>3.  When  the  oxide  of  a  metal  is  heated  with  fused 
borax,  it  dissolves  as  the  result  of  its  union  with  the  excess  of  boric 


438  INORGANIC  CHEMISTRY  FOR  COLLEGES 

anhydride  to  form  a  metaborate.  For  example,  copper  oxide  and 
borax  under  these  conditions  form  a  mixture  of  sodium  metaborate 
and  copper  metaborate,  which  is  fusible: 

CuO  +  Na2O,2B2O3  =  CuO,B2O3  +  Na2O,B2O3 

A  similar  reaction  is  made  use  of  in  welding  together  two  pieces  of 
iron;  the  parts  to  be  united  are  covered  with  borax,  heated  to 
redness,  and  hammered;  the  oxide  on  the  surface  of  the  metal 
does  not  melt  at  the  temperature  used,  and  is  converted  into  iron 
metaborate,  which  mixes  with  the  molten  borax  and  is  thus  expelled 
from  the  point  of  contact  of  the  two  pieces  of  metal;  they  can  now 
be  welded  together  in  the  soft  condition  of  the  iron  at  the  tem- 
perature used.  Borax  is  used  for  the  same  reason  in  soldering. 

The  property  shown  by  borax  of  dissolving  oxides  of  metals 
is  utilized  in  qualitative  chemical  analysis  in  the  so-called  borax- 
bead  test.  A  bit  of  borax  is  placed  in  a  small  loop  in  the  end  of  a 
platinum  wire  and  melted  down  to  a  clear  glass  in  the  Bunsen 
flame.  While  hot  it  is  brought  into  contact  with  the  substance 
to  be  tested  and  returned  to  the  flame.  At  the  high  tempera- 
ture, salts  are  decomposed  and  the  resulting  oxides  dissolve  in  the 
molten  borax  as  metaborates.  Many  of  these  salts  possess  char- 
acteristic colors,  and  the  metal  present  can,  therefore,  be  iden- 
tified. Further,  in  some  cases  when  a  metal  forms  salts  in  which  it 
can  exhibit  different  valencies,  the  color  of  the  salts  is  different 
when  it  is  in  one  condition  from  that  when  in  the  other.  If  the 
bead  is  held  in  the  oxidizing  part  of  the  Bunsen  flame  near  the 
tip,  the  metaborate  formed  will  be  derived  from  the  metal  in  its 
higher  state  of  oxidation;  if  it  is  held  in  the  reducing  flame,  near 
the  top  of  the  blue  inner  cone,  the  salt  will  be  derived  from  the 
metal  in  its  lower  state  of  oxidation. 

522.  Microcosmic  salt,  sodium-ammonium  phosphate  (462), 
can  be  used  instead  of  borax  in  the  bead  test.  This  salt  is 
converted  into  sodium  metaphosphate,  NaPOs,  when  fused, 
and  when  heated  with  oxides  of  metals  dissolves  them  with  the 
formation  of  orthophosphates;  for  example,  with  copper  oxide  the 
following  reaction  takes  place: 

CuO  +  NaP03  =  CuNaP04 


SILICON  AND  BORON  439 

523.  Tests  for  Boric  Acid  and  Borates.— The  most  important 
test  for  these  substances  is  based  on  the  fact  that  certain  organic 
salts  of  boric  acid  impart  a  green  color  to  the  flame  produced  when 
they  burn.     The  test  is  carried  out  by  treating  the  substance  to 
be  examined  with  alcohol  and  concentrated  sulphuric  acid;  when  a 
flame  is  applied  to  the  mixture,  the  alcohol  and  the  ethyl  borate, 
(C2H5)3BO3,  formed  burn,  and  a  green  flame  is  produced. 

A  second  test  which  is  convenient  is  carried  out  by  first  moist- 
ening a  piece  of  turmeric  paper  with  a  solution  of  the  borate  con- 
taining a  little  hydrochloric  acid,  and  then  drying  the  paper  by 
placing  it  on  a  flask  containing  boiling  water.  The  paper,  which 
was  originally  yellow,  turns  to  a  rose  pink  on  drying  if  boric  acid  is 
present,  and  if  put  into  a  dilute  solution  of  ammonia  the  pink  color 
changes  to  bluish  black. 

524.  Other  Compounds  of  Boron. — Boron  forms  two  hydrides, 
B4Hio,  and  B2He,  a  trichloride,  and  a  trifluoride.     Boron  trifluoride 
can  be  prepared  by  a  reaction  analogous  to  that  used  to  make 
silicon  fluoride,  and,  like  the  latter  compound,  reacts  with  water 
to  form  a  complex  acid;  a  number  of  salts  of  hydrofluoboric  acid, 
HBF4,  are  known. 

THE  ACID-FORMING  ELEMENTS  AND  THE  PERIODIC  LAW 

525.  The  elements  described  up  to  this  point  are  the  more 
important  ones  possessing  acid-forming  properties.     From  time  to 
time  the  relation  between  their  physical  and  chemical  properties 
and  their  position  in  the  periodic  classification  of  the  elements 
has  been  emphasized,  and  the  effect  of  the  valence  of  an  element 
on  the  properties  of  the  compounds  containing  it  has  been  pointed 
out.     It  will  be  well  worth  while  to  bring  together  here  the  more 
important  generalizations,  as  these  will  help  materially  to  fix  in 
the  mind  facts  already  learned,  and  will  be  the  best  preparation  for 
the  study  of  the  chemistry  of  the  metals  and  their  compounds. 

526.  Relationships  in  a  Chemical  Family. — Physical  Proper- 
ties.    With  increasing  atomic  weight  the  densities  of  the  elements 
and  their  compounds  increase,  and  their  boiling-points,  melting- 
points,  and  critical  temperatures  also  increase.     The  number  of  atoms 
in  the  molecules  of  the  elements  in  the  gaseous  condition,  either 
remains  the  same,  as  in  the  case  of  the  halogens,  or  decreases  as 
seen  in  the  phosphorus  family. 


440  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Chemical  Properties.  The  stability  toward  heat  of  the  molecule 
of  the  element  in  the  gaseous  condition  decreases  with  increasing 
atomic  weight;  this  is  shown  in  the  cases  of  chlorine  and  iodine 
and  phosphorus  and  arsenic. 

The  activity  of  the  elements  as  measured  by  the  heat  of  forma- 
tion of  the  hydrides  and  other  compounds  containing  electro- 
positive elements  decreases  with  increasing  atomic  weight,  a  fact 
which  is  associated  with  the  decreasing  stability  of  the  hydrides 
with  increasing  atomic  weights.  The  heats  of  formation  of  the 
hydrides  from  the  elements  in  the  gaseous  condition  are  positive  in 
most  cases.  As  the  affinity  of  the  elements  for  hydrogen  decreases 
their  electro-positive  properties  increase.  Since  in  the  salts  of  the 
acids  containing  oxygen  the  element  is  functioning  in  its  positive 
character,  the  element  with  the  higher  atomic  weight  will  replace, 
in  general,  the  one  with  the  lower  atomic  weight.  On  the  other 
hand,  in  the  case  of  the  hydrides,  in  which  the  elements  are 
electro-negative,  the  order  of  replacement  is  the  reverse;  chlorine, 
for  example,  replaces  bromine. 

The  extent  of  the  hydrolysis  of  the  chlorides  of  the  elements 
decreases  with  increasing  atomic  weight,  which  is  another  indica- 
tion of  increase  in  electro-positive  nature  in  this  order. 

The  effect  of  change  in  valence  of  an  element  is  great.     Most  of 
the  oxides  of  the  acid-forming  elements  unite  with  water  to  form 
acids,  the  strength  of  which  in  the  case  of  a  single  element  is 
determined  by  its  valence — the  higher  the  valence,  that  is,  the 
larger  the  proportion  of  oxygen,  the  stronger  the  acid.     We  have 
seen,  for  example,  that  nitric,  sulphuric,  and  chloric  acids  are 
stronger,  respectively,  than  nitrous,  sulphurous,  and  hypochlorous 
acids.     In  the  case  of  amphoteric  elements  like  antimony,  the 
hydroxide  containing  the  element  in  the  lower  valence  is  a  weaker 
acid  and  a  stronger  base  than  the  hydroxide  containing  the  ele- 
ment in  the  higher  valence.     We  shall  see  later  that  this  important 
generalization  can  be  extended  to  include  the  bases  formed  from 
the  metallic  elements;   the  lower  the  valence  of  the  element  the 
stronger  the  base  it  forms;    sodium  hydroxide,  NaOH,  is  a  very 
active  base,  aluminium  hydroxide,  A1(OH)3,  is  a  very  weak  base, 
which  shows  some  of  the  properties  of  acids;  iron  forms  two  basic 
hydroxides,  Fe(OH)2  and  Fe(OH)3,  the  former  being  the  stronger 
base. 


SILICON  AND  BORON  441 

527.  Relationships  between  the  Chemical  Families. — When 
the  elements  were  taken  up  systematically  according  to  the  periodic 
classification,  the  halogen  family  was  first  discussed  because  the 
electro-negative  character  of  all  its  members  is  pronounced,  and 
fluorine  is  the  most  electro-negative  of  all  the  elements.  As  we 
consider  the  families  one  after  another  passing  from  the  right  of 
the  table  to  the  left,  we  discover  that  the  compounds  of  the  ele- 
ments with  oxygen  and  hydrogen  become  weaker  and  weaker 
acids,  and  the  electro-negative  character  of  the  elements  decreases. 
This  is,  no  doubt,  associated  with  their  decreasing  valence  toward 
oxygen.  Since  in  any  family  decrease  in  electro-negative  nature 
follows  increase  in  atomic  weight,  we  have,  thus,  two  factors  which 
tend  simultaneously  to  decrease  the  electro-negative  character  of 
the  elements.  While  in  the  seventh  group  all  the  elements  are 
strongly  electro-negative,  in  the  sixth  tellurium  shows  some 
metallic  properties,  in  the  fifth  antimony  and  bismuth,  in  the  fourth 
all  the  members  below  silicon,  and  in  the  third  even  boron,  the 
first  member  of  the  group,  shows  these  properties,  although  its 
oxide  is  more  acidic  than  basic.  If  we  draw  a  line  across  the 
periodic  table  from  the  square  in  which  boron  is  placed  to  that 
under  iodine,  we  divide  the  elements  roughly  into  those  which  are 
electro-positive  and  those  which  are  electro-negative;  the  ele- 
ments near  the  diagonal  line  show  amphoteric  properties;  the 
most  electro-positive  and  electro-negative  elements  are  furthest 
from  the  line.  The  classification  in  this  way  does  not  include  the 
elements  in  group  0  which  show  no  chemical  properties,  and  those 
in  group  8,  which  are  distinctly  metallic  in  properties. 

EXERCISES 

1.  Write  balanced  equations  for  the  reactions  which  take  place  between 
the  following  substances:    (a)  silicon  and  aqua  regia,    (6)  water-glass  and 
hydrochloric  acid,   (c)  Colemanite  and  sodium  carbonate  when  fused  together, 
(d)  silicon   tetrachloride,    water,    and   ammonia,     (e)  silicon   dioxide,    con- 
centrated sulphuric  acid,  and  calcium  fluoride,    (/)  silicon  dioxide  and  calcium 
sulphate   (fused),     (g)  lime,   sodium  sulphate,  and  silicon  dioxide   (fused), 
(h)  lead  oxide  and  water-glass  (fused),     (i)  Colemanite,  water,  and  sulphur 
dioxide,    (j)  boron  nitride  and  steam,    (fc)  boracite  and  hydrochloric  acid, 
(Z)  borax  and  hydrochloric   acid,     (ra)  boric   acid  and   sodium   hydroxide, 
(n)  boric  acid,  alcohol  (C2H5OH),  and  concentrated  sulphuric  acid. 

2.  Give  as  many  examples  as  you  can  of  the  facts  summarized  in  section 
526.     Tabulate  the  facts  according  to  the  properties  listed. 


CHAPTER    XXXI 
PHYSICAL  PROPERTIES  OF  METALS.     ALLOYS 

The  uses  to  which  metals  are  put  are  based  upon  their  physical 
or  chemical  properties,  but  as  the  great  majority  of  these  uses 
depend  on  physical  properties  the  latter  will  be  discussed  briefly, 
although  their  consideration  rightly  falls  in  physics  and  engi- 
neering. 

528.  Density. — The   metals   vary   greatly   in    density    (175); 
the  lightest  is  lithium,  which  has  the  density  0.534  and  is,  there- 
fore, about  one-half  as  heavy  as  water,  and  the  heaviest  is  osmium 
(d.  22.5)  which  is  closely  related  to  platinum  (d.  21.37)  in  physical 
and  chemical  properties.     The  so-called  light  metals,  of  which 
sodium,   potassium,   magnesium,   and   aluminium   are   examples, 
have  a  density  less  than  4;   iron,  lead,  tin,  silver,  etc.,  are  known 
as  heavy  metals.     In  the  construction  of  automobiles  and  aero- 
planes the  density  of  the  metal  used  is  an  important  factor,  for  the 
power  consumed  in  moving  them  increases  rapidly  as  the  weight 
increases.     For  this  reason  aluminium  (d.  2.7)  is  used  for  such 
parts  as  do  not  bear  a  great  strain;  the  metal  cannot  replace  the 
much  heavier  steel  for  many  purposes  on  account  of  the  fact  that 
its  tensile  strength  (see  below)  is  comparatively  low.     An  alloy 
which  contains  aluminium,  copper,  and  magnesium  is  used  in 
making  parts  for  aeroplanes.    Magnesium  (d.  1.74)  is  not  a  strong 
metal,  but  forms  alloys  with  other  metals  which  have  a  low  density 
and  a  relatively  high  tensile  strength.     The  table  on  page  443  gives 
the  densities  and  some  other  physical  properties  of  the  metals ;   it 
should  be  consulted  in  connection  with  the  discussion  of  these 
properties  which  follows. 

529.  Hardness. — The  metals  vary  in  hardness,  from  potas- 
sium, which  can  be  molded  like  wax,  to  chromium,  which  will  cut 
glass.     In  the  arbitrary  scale  of  hardness  which  is  used  to  express 
numerically  this  property,  a  number  of  substances  are  selected  as 

442 


PHYSICAL  PROPERTIES  OF  METALS.     ALLOYS 


443 


standards  (see  footnote  4,  below)  and  arranged  in  such  an  order 
that  each  member  of  the  series  can  be  scratched  by  the  succeeding 
member.  When  the  hardness  of  lead  is  said  to  be  1.5  the  state- 
ment means  that  lead  will  scratch  talc,  1,  but  not  rock-salt,  2. 


PHYSICAL  PROPERTIES  OF  METALS 


Metal. 

Density. 

Melting- 
point. 

Boiling- 
point. 

Electrical  l 
Eesistance. 

Heat  2 
Conduc- 
tivity. 

Tensile  3 
Strength. 

Hard- 
ness.4 

Aluminium  . 
Antimony.  . 
Barium 

2.70(20°) 
6.62(20°) 
380 

658° 
630° 
850° 

1800° 
1440° 
950° 

2.6-3.0 
35.4-45.8 

0.48 
0.0444 

12,590 
1,000 

2. 
3.3 

Bismuth  
Cadmium 

9.78(20°) 
8  65(20°) 

269° 
321° 

1420° 

778° 

108 
6.2-7.0 

0.017 
0.22 

3,000 

2.5 

Caesium 

1  87(26°) 

264° 

670° 

193 

Calcium 

1.54(29°) 

805° 

10.5 

Copper  
Gold  
Iron  
Lead  
Lithium 

8.94 
19.32(17°) 
7.90 
11.34 
0534 

1083° 
1063° 
1520° 
327° 
186° 

2310° 
2200° 
2450° 
1525° 
1400° 

1.54 
2.09 
9.7-12.0 
18.4-19.6 

85 

1.00 
0.75 
0.20 
0.084 

24,000 
20,000 
48,000 
2,050 

2.5-3 
2.5-3 
4-5 
1.5 

]VIagnesium 

1  74(5°) 

651° 

1120° 

4  1-50 

037 

Mercury  .  .  . 
Nickel  

13.595(4°) 
8.8 

-38.85° 
1452° 

357.25° 
2450° 

94 
10.7-12.4 

0.015 
0.142 

54,000 

Platinum  .  .  . 
Potassium 

21.37 
0  87(20°) 

1755° 
625° 

2650° 
712° 

9.0-15.5 

0.166 

45,000 

4.3 

Silver  
Sodium 

10.50 
097 

961° 
976° 

1955° 
750° 

1.5-1.7 
4.4 

1.09 

41,000 

2.5-3 

Strontium  . 

2.54 

900° 

24.8 

Tin 

7.3(15°) 

231.9° 

2270° 

9.53-11.4 

0.153 

4,600 

1.5 

Tungsten 

19  1 

3200° 

70 

Zinc      .    . 

7.19 

419° 

940° 

5.56-6.04 

0.265 

5,000 

2.5 

1  The  figures  given  are  expressed  in  millionths  of  an  ohm  and  are  the 
resistance  offered  by  a  cube  of  the  metal  1  centimeter  on  each  edge. 

2  The  figures  given  express  in  calories  the  quantity  of  heat  transmitted 
per  second  through  a  plate  1  centimeter  thick  across  an  area  of  1  square 
centimeter  when  the  difference  in  temperature  between  the  two  sides  of  the 
plate  is  one  degree. 

3  The  figures  give  approximate  values  of  the  number  of  pounds  per  square 
inch  required  to  effect  a  permanent  elongation  of  the  cast  metal. 

4  The  arbitrary  scale  of  hardness  used  is  as  follows :    1  talc,  2  rocksalt, 
3  calcite,  4  fluorite,  5  apatite,  6  feldspar,  7  quartz,  8  topaz,  9  corundum, 
10  diamond, 


444  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  hardness  of  metals  is  markedly  affected  by  the  presence  of 
other  substances  in  them.  While  the  hardness  of  pure  iron  is 
4  to  5,  steel,  which  contains  carbon,  silicon,  and  other  substances, 
varies  in  hardness  according  to  its  composition  from  5  to  8.5; 
the  use  of  steel  in  making  files  is  based  on  its  hardness;  sandpaper, 
which  is  made  by  covering  paper  with  quartz,  hardness  7,  is  used 
for  polishing  wood  and  other  soft  objects,  but  emery  paper,  made 
from  crushed  corundum,  hardness  9,  is  used  when  iron  and  other 
metals  are  to  be  abraded.  The  position  of  a  few  common  sub- 
stances in  the  scale  of  hardness  is  of  interest;  amber  is  2.5,  asphalt, 
1-2,  brass  3-4,  garnet  7,  glass  4.5-6.5,  marble  3-4,  meerschaum  2-3. 

530.  Tenacity. — The    metals  and  other  substances  differ  in 
the  extent  to  which  they  can  resist  a  strain  brought  to  bear  on  them 
that  tends  to  bring  about  a  permanent  change  in  their  form.     In 
the  seventh  column  of  the  table  on  page  443  are  given  the  values 
of  the  tensile  strength  of  the  more  common  metals.     This  property 
is  greatly  affected  by  the  purity  of  the  metal  and  the  physical 
treatment  to  which  it  has  been  subjected,  such  as  casting,  rolling, 
drawing,  etc.     Pure  iron  wire,  hard  drawn,  has  a  tensile  strength 
of  from  80,000  to  120,000;   after  annealing  the  value  is  50,000  to 
60,000.     The  tensile  strength  of  cast  iron  is  from  13,000  to  33,000 
and  of  steel  from  80,000  to  460,000.     Cast  zinc  has  a  tensile 
strength  of  from  7,000  to  13,000,  whereas  the  value  for  the  drawn 
metal  is  22,000  to  30,000.     The  effect  of  heat  treatment  on  the 
properties  of  steel  will  be  considered  later  in  more  detail  on  account 
of  its  great  importance. 

531.  Electrical  Conductivity. — All  substances  offer  more  or  less 
resistance  to  the  flow  of  an  electric  current  through  them.     With 
any  given  substance  the  resistance  is  determined  by  its  dimen- 
sions and  the  temperature.     When  these  are  fixed  the  resistance 
is  constant  and  a  characteristic  property  of  the  substance.     Differ- 
ent substances  differ  widely  in  the  resistance  they  offer  to  the  flow 
of  an  electric  current;  in  the  case  of  metals  it  is  very  small,  whereas 
with  the  non-metals  and  most  other  substances  in  the  solid  condi- 
tion it  is  very  great.     The  standard  of  resistance  is  that  offered 
at  0°  by  a  column  of  mercury  1  mm.  in  cross-section  and  1.063 
meters  in  length;  it  is  called  an  ohm. 

Electrical  conductivity  is  the  reverse  of  electrical  resistance 
and  is  inversely  proportional  to  it.     For  example,  the  resistance 


PHYSICAL  PROPERTIES  OF  METALS.    ALLOYS  445 

of  silver  is  about  one-sixtieth  that  of  mercury,  and,  consequently, 
its  conductivity  is  sixty  times  that  of  the  latter.  The  electrical 
conductivity  is  denned  as  the  reciprocal  of  the  resistance;  it  equals 
numerically  1  divided  by  the  resistance,  and  is  expressed,  there- 
fore, in  reciprocal  ohms. 

In  the  fifth  column  in  the  table  on  page  443  the  specific  resist- 
ance of  the  metals  is  given,  that  is,  the  resistance  of  a  cube  of 
the  metal  1  cm.  on  each  edge.  Since  the  numbers  are  very  small 
they  are  expressed  in  millionths  of  an  ohm  to  avoid  large  fractions. 
It  is  seen  that  there  is  a  rather  wide  variation  in  the  values  for 
the  several  metals;  the  resistance  of  silver,  copper,  gold,  and 
aluminium  is  comparatively  small;  magnesium  and  zinc  come  next, 
and  platinum,  iron,  and  nickel  fall  into  a  group.  The  resistance  of 
the  elements  in  which  metallic  properties  are  not  highly  developed, 
like  antimony  and  bismuth,  is  high. 

The  conductivity  of  a  metal  is  affected  by  its  physical  condi- 
tion brought  about  as  the  result  of  heat  treatment,  drawing, 
casting,  etc.,  and  by  the  presence  of  even  very  small  amounts  of 
impurities.  It  is  for  this  reason  that  copper  is  refined  with  the 
greatest  care  when  it  is  to  be  employed  in  the  form  of  wire  for  elec- 
trical construction  on  account  of  its  low  resistance.  The  figures 
given  in  the  table  for  iron,  9.7-12.0,  indicate  the  resistance  of  the 
usual  grades  of  commercial  iron  wire;  using  the  same  unit  the  re- 
sistance of  hard  cast  iron  is  97.8,  and  that  of  a  soft  steel  15.9 
and  a  hard  steel  45.7.  The  high  resistance  of  certain  alloys  is 
utilized  in  the  construction  of  resistance  coils  for  electrical 
purposes  and  when  it  is  desired  to  convert  electrical  energy 
into  heat  to  be  used  in  connection  with  electric  stoves,  ovens, 
and  furnaces.  The  heat  generated  by  an  electric  current  is 
proportional  to  the  resistance.  The  alloy  known  as  constantan, 
which  contains  60  per  cent  of  copper  and  40  per  cent  of  nickel,  is 
used  for  this  purpose;  its  specific  resistance  in  the  unit  used  in  the 
table  is  about  49.  Platinum  can  also  be  used  and  is  of  particular 
value  because  it  does  not  oxidize  at  high  temperatures,  but  its 
Cost  prohibits  its  use  except  in  small  apparatus.  Nichrome  (or 
chromel)  is  an  alloy  of  nickel  and  chromium  which  is  much  used  in 
the  construction  of  electrical  heating  devices.  Its  specific  resist- 
ance (about  100)  is  great  and  it  resists  the  action  of  air  at  reason- 
ably high  temperatures. 


446  INORGANIC  CHEMISTRY  FOR  COLLEGES 

It  is  of  interest  to  compare  the  resistance  of  metals  with  that 
of  the  so-called  "  non-conductors  "  or  dielectrics.  The  resistance 
of  a  centimeter  cube  of  copper  is  1.5  millionths  of  an  ohm,  while 
that  of  glass  of  the  same  dimensions  is  9  X  1013  ohms.  The  values 
for  a  few  other  substances  are  as  follows:  quartz  crystal  1  X  1014, 
rock  salt  9  X  1016,  rhombic  sulphur  70  to  390  X  1013,  dry  wood 
5  to  10  X  108,  petroleum  2  X  1016,  and  distilled  water  0.5  X  106. 

Solutions  of  acids,  bases,  and  salts  conduct  electricity.  The 
following  figures  express  in  ohms  the  resistance  of  10  per  cent 
solutions:  hydrochloric  acid  1.59,  sodium  hydroxide  3.20,  sodium 
chloride  8.33,  zinc  chloride  13.75,  and  copper  sulphate  31.2. 

532.  Heat  Conductivity. — Metals  are  the  best  conductors  of 
heat.     In  the  sixth  column  of  the  table  of  the  physical  properties 
of  metals  will  be  found  the  quantity  of  heat,  measured  in  calories, 
transmitted  per  second  through  a  centimeter  cube  of  the  metal 
when  the  difference  in  temperature  of  the  two  sides  of  the  cube  is 
1  degree.     A  comparison  of  the  figures  with  those  for  the  electrical 
resistance  of  the  metals  brings  out  clearly  the  fact  that  high  heat 
conductivity  is  associated  with  low  electrical  resistance;  the  good 
conductors  for   electricity   conduct   heat  well;    high   heat    con- 
ductivity  is,    thus,    a    characteristic    property   of   metals.     The 
corresponding  figures  for  some  non-conductors  of  heat  are  as  fol- 
lows:   asbestos  paper,  which  is  used  in  making  fire-proof  paper, 
theater  curtains,  etc.,  0.00043  calorie;  red  brick  0.00150;  blotting 
paper  0.00015;   cork  0.00072;   cotton  wool,  0.000043;   eiderdown, 
0.000011;   felt,  0.000087;   glass,  0.0012;    ice,  0.00396;   magnesia, 
0.00016-0.00045;      quartz,     0.00036;      sawdust,     0.00012;     silk, 
0.00009;   dry  soil,  0.00033;  air,  0.000057;  water,  0.0012. 

533.  The  Physical  State  of  Metals. — All  the  metals  with  the 
exception  of  mercury  are  solids  at  ordinary  temperatures.     They 
can  be  obtained  in  the  form  of  crystals  by  fusing  them  and  pouring 
off  the  molten  metal  after  a  part  has  crystallized;  or  they  can  be 
crystallized  from  a  solvent,  using  for  this  purpose  some  other 
metal  with  which  they  do  not  interact.     Most   of  the  metals 
crystallize   in  the  cubic  system.     When  molten  metals  solidify 
they  change  to  a  mass  of  closely  interlocked  crystals;    if  such 
a  mass  is  hammered,  rolled,  or  drawn  through  a  die,  the  crystals 
are  more  or  less  broken,  and  due  to  the  pressure  exerted  on  them 
their  faces  are  forced  into  closer  contact  and,  perhaps,  converted 


PHYSICAL  PROPERTIES  OF  METALS.    ALLOYS          447 

into  an  amorphous  condition  when  they  coalesce.  If  a  piece  of 
metal,  such  as  a  wire,  is  heated  until  it  softens,  the  crystals  have  an 
opportunity  to  grow  and  rearrange  themselves;  if  now  it  is  cooled 
rapidly  by  plunging  it  into  water,  it  is  "  frozen  "  in  the  condition 
in  which  it  existed  at  the  higher  temperature,  and  its  physical 
properties  are  different  from  what  they  were  before  the  heat 
treatment.  If  a  platinum  wire  which  has  become  brittle  is  heated 
to  redness  and  plunged  into  water  it  becomes  soft  and  pliable. 
Most  metals  are  hardened  by  this  treatment.  It  is  evident  that 
pressure  and  heat  treatment  have  an .  effect  on  the  physical  prop- 
erties of  metals;  this  is  more  notably  the  case  with  alloys  in  which 
there  are  several  kinds  of  crystals  present  in  the  solid  material. 

The  color  of  metals  is  determined  by  their  state  of  division. 
The  color  usually  associated  with  a  metal  is  that  of  the  light 
reflected  from  its  surfaces  when  it  is  in  the  massive  condition; 
silver  is  white,  gold  yellow,  copper  red,  etc.  In  very  thin  layers 
metals  transmit  light,  which  in  certain  cases  differs  in  color  from 
that  reflected  from  the  surface;  gold  and  copper,  for  example, 
are  green  when  seen  by  transmitted  light. 

534.  Metals  can  be  obtained  in  a  state  of  fine  division  and 
their  color  is  determined  by  the  size  of  the  particles;  metallic  silver, 
for  example,  can  be  obtained  by  precipitation  from  solutions  of  its 
salts  in  forms  which  are  red,  green,  yellow,  etc.  A  convenient  way 
to  obtain  metals  in  the  finely  divided  condition  is  to  produce 
sparks  under  water  by  alternately  bringing  into  contact  and  sep- 
arating the  terminals  of  an  electric  circuit.  If  the  terminals  are 
made  of  gold,  each  time  they  are  separated  and  a  spark  is  pro- 
duced, some  of  the  metal  is  torn  away  in  the  form  of  particles  so 
small  that  they  stay  suspended  in  the  water  and  cannot  be  seen 
as  such  by  the  eye;  in  this  condition  gold  is  bright  red  in  color  and 
appears  to  be  in  solution.  It  is  in  what  is  known  as  the  colloidal 
condition.  When  many  substances  insoluble  in  water  are  in  an 
exceedingly  finely  divided  condition  they  appear  to  dissolve  and 
cannot  be  removed  by  filtration.  Platinum  in  the  form  of  a  wire 
or  sheet  is  a  silver-white  metal;  in  the  colloidal  condition  it  is 
black.  Metals  when  in  the  finely  divided  condition  exhibit  prop- 
erties which  are,  no  doubt,  due  to  the  fact  that  the  surface  of  the 
metal  is  almost  infinitely  greater  than  when  it  is  in  the  massive 
state.  Iron,  for  example,  reacts  but  slowly  with  the  oxygen  of 


448  INORGANIC  CHEMISTRY  FOR  COLLEGES 

the  air  when  it  is  in  the  usual  condition;  it  can  be  obtained  so 
finely  divided  that  it  takes  fire  spontaneously  when  brought  into 
the  air.  Finely  divided  metals  serve  as  valuable  catalytic  agents; 
platinum  is  used  to  hasten  the  union  of  sulphur  dioxide  and  oxygen 
in  the  manufacture  of  sulphuric  acid,  and  nickel  in  the  preparation 
of  solid  cooking  fats  by  the  addition  of  hydrogen  to  vegetable  oils. 

535.  Change  in  State  of  the  Metals.— All  the  metals  melt  and 
can  be  vaporized,  although  in  certain  cases  a  temperature  attain- 
able only  in  an  electric  furnace  is  necessary.  The  melting-points 
and  boiling-points  of  the  commoner  metals  are  given  in  the  table 
on  page  443.  The  range  in  melting-points  is  great;  mercury  melts 
at  —38.8°,  caesium,  a  metal  that  resembles  sodium  in  properties, 
at  26.4°  and  tungsten  at  3200°.  The  use  of  mercury  in  thermom- 
eters and  that  of  tungsten  in  electric  lamps  are  based  largely  on 
the  melting-points  of  the  metals.  Metals  that  boil  below  1000° 
can  be  more  or  less  readily  purified  by  distillation.  The  fact  that 
mercury  boils  at  a  much  lower  temperature  than  any  other  metal 
makes  it  possible  to  obtain  it  in  a  very  pure  condition  for  use  in 
such  measuring  instruments  as  thermometers,  barometers,  and 
apparatus  employed  in  the  analysis  of  gases.  Many  metals  dissolve 
in  mercury  and  their  presence  even  in  small  amounts  makes  it 
cling  to  glass  and  leave  a  film  on  its  surface;  it  must,  accordingly,  be 
separated  from  these  before  it  is  used  for  the  above  purposes.  The 
fact  that  mercury  dissolves  other  metals  and  is  readily  volatilized 
is  used  in  one  process  of  extracting  gold  from  its  ores.  The 
crushed  ore  is  stirred  with  mercury  and  the  solution  of  gold  thus 
formed  is  then  heated  to  the  temperature  at  which  mercury  boils; 
after  the  latter  has  distilled  off  the  residue  is  gold. 

When  molten  metals  change  from  the  liquid  to  the  solid  con- 
dition there  is  a  change  in  volume;  in  most  cases  contraction  takes 
place.  Gold  cannot  be  cast,  because  in  passing  from  the  liquid 
to  the  solid  condition  it  contracts  and,  therefore,  shrinks  on  solidi- 
fication from  the  mold  into  which  it  is  poured.  Since  a  sharp 
impression  cannot  be  obtained  in  this  way,  in  making  coins,  gold 
is  struck  with  a  die  while  hot.  Bismuth  expands  on  solidification. 
Antimony  is  put  into  lead  which  is  to  be  used  to  cast  type,  because 
it  forms  with  the  metal  an  alloy  that  expands  when  it  solidifies. 

536.  Vapor-density  of  the  Metals. — A  few  of  the  metals  boil 
at  a  sufficiently  low  temperature  to  make  it  possible  to  determine 


PHYSICAL  PROPERTIES  OF  METALS.    ALLOYS  449 

their  vapor  densities;  in  all  cases  that  have  been  examined  the 
results  lead  to  the  conclusion  that  the  elements  that  are  distinctly 
metallic  in  character,  like  sodium  and  mercury,  are  monatomic 
in  the  gaseous  state.  Antimony,  which  is  more  non-metallic 
than  metallic  is  diatomic,  and  bismuth,  which  more  nearly  ap- 
proaches a  metal  in  physical  and  chemical  properties,  yields  a 
mixture  of  monatomic  and  diatomic  molecules  at  a  temperature 
just  above  its  boiling-point, 


ALLOYS 

637.  The  solids  obtained  when  two  or  more  metals  are  mixed 
in  the  molten  condition  and  allowed  to  solidify  are  called  alloys. 
On  account  of  the  fact  that  the  presence  of  certain  non-metallic 
elements,  such  as  phosphorus  and  antimony,  have  a  marked  effect 
on  the  properties  of  the  materials  made  in  this  way,  these  ele- 
ments are  frequently  added  to  the  metals  in  the  preparation  of 
alloys.  The  so-called  non-ferrous  alloys  only  will  be  considered 
briefly  here;  those  which  contain  iron,  such  as  the  various  kinds 
of  steel,  are  of  such  technical  importance  as  to  demand  separate 
treatment  when  the  chemistry  of  iron  is  discussed. 

The  conditions  in  regard  to  solubility  which  exist  when  molten 
metals  are  mixed  are  the  same  as  those  obtained  when  substances 
liquid  at  the  ordinary  temperatures  are  brought  together.  It  will 
be  recalled  that  liquids  either  mix  in  all  proportions,  exhibit  definite 
solubilities  one  in  the  other,  or  are  practically  immiscible.  Exam- 
ples of  these  classes  are,  respectively,  alcohol  and  water,  ether  and 
water,  and  kerosene  and  water.  Similar  relationships  are  found 
among  the  metals  in  the  liquid  state;  examples  of  the  three  classes 
just  mentioned  are,  respectively,  silver  and  copper,  lead  and  tin, 
and  lead  and  copper.  In  the  case  when  the  metals  exhibit  a 
definite  solubility,  one  in  the  other,  the  solubilities,  as  is  the  case 
in  general  with  all  types  of  solutions,  vary  with  the  temperature. 
All  these  relationships  come  into  play  when  molten  mixtures 
of  metals  cool  and  finally  solidify,  and  they  determine  the  physical 
structure  of  the  solid  alloy. 

In  the  case  of  the  first  class  of  alloys,  namely,  those  made  up 
of  metals  miscible  in  all  proportions,  the  liquid  mixture  solidifies 


450  INORGANIC  CHEMISTRY  FOR  COLLEGES 

to  a  solid  uniform  in  appearance  when  examined  with  a  microscope 
of  the  highest  power.  The  metals,  apparently,  are  as  evenly  dis- 
tributed in  the  solid  condition  as  they  were  in  the  liquid  state; 
they  are  said  to  exist,  accordingly,  in  solid  solution.  The  melting- 
points  of  such  alloys  change  gradually  as  the  composition  of  the 
mixture  is  changed  and  are  determined  by  the  relative  amounts  of 
the  metals  present.  Alloys  of  the  third  type  are  also  simple  in 
structure.  In  this  case  the  immiscible  mixture  may  be  stirred 
to  bring  the  constituents  into  more  intimate  contact  and  then 
allowed  to  solidify.  On  microscopic  examination  they  show  two 
distinct  kinds  of  solid  particles  which  are  made  up  of  the  two 
metals  present.  When  such  an  alloy  is  gradually  heated  to  deter- 
mine its  melting-point  the  lower  melting  metal  first  liquefies  at  its 
melting-point,  and  then  the  thermometer  remains  constant  for  a 
time  as  the  heat  supplied  is  used  in  melting  the  more  fusible  metal ; 
after  this  has  taken  place  the  thermometer  rises  until  the  melting- 
point  of  the  second  metal  is  reached,  when  it  again  comes  to  rest. 
Such  alloys  show,  therefore,  two  distinct  melting-points,  which  are 
those  of  its  constituents. 

538.  The  conditions  that  exist  in  the  class  of  alloys  made  up  of 
metals  that  show  limited  solubilities  in  each  other  are  much  more 
complicated  and  yield  products  more  complex  in  structure.  The 
alloys  made  from  lead  and  tin,  which  are  used  as  solder,  are  exam- 
ples of  this  class.  Lead  melts  at  327°,  and  tin  at  232°.  When 
tin  is  added  to  molten  lead,  the  temperature  at  which  the  mixture 
begins  to  solidify  is  lower  than  the  melting-point  of  lead,  and  like- 
wise a  mixture  of  tin  and  a  small  amount  of  lead  begins  to  solidify 
at  a  temperature  below  the  melting-point  of  tin.  In  both  cases 
the  melting-point  is  lowered — a  fact  in  accord  with  the  generaliza- 
tion already  given,  namely,  that  the  freezing-point  of  a  solvent  is 
lowered  by  the  presence  of  a  dissolved  substance.  In  the  accom- 
panying diagram  (Fig.  32)  are  plotted  the  temperatures  at  which 
liquid  mixtures  of  lead  and  tin  begin  to  solidify. 

The  significance  of  the  diagram  will  be  appreciated  from  the 
following  discussion:  The  point  marked  x  in  the  diagram  repre- 
sents a  mixture  made  up  of  70  per  cent  of  lead  and  30  per  cent  of 
tin  at  the  temperature  300°;  at  this  temperature  the  mixture  is 
liquid.  If  the  temperature  falls,  we  see  by  following  down  the 
vertical  line  that  at  about  250°  solidification  begins;  at  this  point 


PHYSICAL  PROPERTIES  OF  METALS.    ALLOYS 


451 


lead  begins  to  separate  in  the  solid,  crystalline  condition,1  and  this 
separation  continues  with  falling  temperature  until  180°  is  reached, 
when  the  entire  mass  becomes  solid.  If  we  examine  next  the 
behavior  of  a  mixture  represented  by  the  position  y  on  the  dia- 
gram— one  containing  80  per  cent  tin  and  20  per  cent  lead  at  250° 
— we  see  that  in  this  case  tin  2  separates  first  at  190°  and  that  the 
entire  mixture  is  solid  at  180°.  The  mixture  having  the  com- 
position indicated  by  the  position  of  B  on  the  diagram — 31  per 
cent  lead  and  69  per  cent  tin — solidifies  completely  at  one  tem- 
perature— it  has  a  definite  melting-point.  The  mixture  which 
has  these  properties  is  called  the  eutectic,  the  name  being  derived 
from  the  Greek  words  meaning  "  well-melting." 


f 


6bO 

A 

300 
250 

X 

.••")(. 

Tin 

\ 

^ 

"^~^Q. 

f^ 

y 

c 

200 
180 

150 
X 

-^^5 

*-**^ 

5 



^^ 

180  Deg. 

D 

1                           •' 

^* 

i  

E 

)           10          20          30          40          50          60          70          &0          90          100 
Per  Cent  Tin 
)0        90         80         70         60         50         40          30         EG          10          0 
Per  Cent  Lead 

FIG.  32. 

The  points  in  the  area  in  the  diagram  above  the  curves  AB 
and  BC  represent  the  composition  and  temperature  of  liquid  mix- 
tures of  the  two  metals;  if  these  conditions  place  a  mixture  in  the 
area  between  the  curves  AB  and  BD,  it  is  composed  of  solid  lead 
and  the  liquid  mixture;  and  if  in  the  area  between  CB  and  BE,  of 
solid  tin  and  the  liquid  mixture;  below  the  line  DE  the  alloy  is  a 
solid. 

It  is  evident  from  the  above  that  an  alloy  made  from  a  mixture 
of  20  parts  of  tin  and  80  parts  of  lead  will  differ  in  properties  from 
one  containing  80  parts  of  tin  and  20  of  lead.  The  former  will 
consist  of  crystals  of  lead  embedded  in  the  eutectic  and  the  latter 
of  crystals  of  tin  in  the  eutectic.  The  physical  structure  of  an 


crystals  formed  are  a  solid  solution  of  about  4  parts  of  tin  in  96 
parts  of  lead. 

2  The  crystals  contain  98  parts  of  tin  and  2  parts  of  lead. 


452  INORGANIC  CHEMISTRY  FOR  COLLEGES 

alloy  is  rendered  visible  by  polishing  a  flat  surface  of  it  and  then 
treating  the  latter  with  a  liquid  which  attacks  it;  if  two  or  more 
substances  are  present  they  are  affected  differently,  and  after  a 
short  time  the  surface  is  found  to  be  etched  in  such  a  way  that  the 
several  ingredients  of  the  alloy  can  be  readily  distinguished  under  a 
microscope.  Figs.  33a,  336,  and  33c  are  reproductions  of  photo- 
micrographs of  alloys  of  tin  and  bismuth,  which  resemble  in 
structure  the  alloys  of  lead  and  tin.  Photographs  of  the  former  are 
used  here  to  illustrate  the  microstructure  of  alloys  because  they 
can  be  reproduced  more  clearly  than  the  photomicrographs  of  the 
lead-tin  alloys.  It  is  evident  that  such  differences  in  physical 
structure  will  have  a  marked  effect  on  the  hardness,  tensile 
strength,  and  other  physical  properties  of  the  alloy. 

Returning  to  the  consideration  of  the  alloys  of  tin  and  lead, 
it  will  be  seen  from  the  diagram  that  if  the  one  described,  rich  in 
lead  (80  lead  to  20  tin),  is  melted  and  allowed  to  cool,  it  begins  to 
solidify  at  about  270°.  It  will  persist  as  a  thick,  viscous  mixture 
made  up  of  solid  and  liquid  until  the  temperature  falls  to  180°, 
when  the  mass  solidifies  completely;  it  will  be  in  a  plastic  condi- 
tion over  a  range  in  temperature  of  90  degrees,  so  that  it  can  be 
worked  or  molded,  as  it  has  about  the  same  viscosity  as  that  of 
baker's  dough.  The  alloy  containing  80  parts  of  tin  and  20  parts 
of  lead  stays  liquid  until  a  lower  temperature  is  reached — about 
190°  instead  of  270°,  the  temperature  at  which  the  alloy  high  in 
lead  begins  to  freeze — but  is  in  the  plastic  condition  during  a  drop 
in  temperature  of  only  10  degrees,  for  the  eutectic  solidifies  at 
180°.  Such  an  alloy  could  be  worked  with  at  low  temperatures, 
but  if  it  is  to  be  handled  in  the  plastic  condition,  as  in  the  case  in 
soldering,  it  would  solidify  so  rapidly  that  the  manipulation  would 
be  very  difficult. 

539.  The  alloys  of  lead  and  tin  are  used  as  solders  and  we  can 
now  understand  why  mixtures  of  different  compositions  are  used 
for  different  purposes.  In  soldering,  two  pieces  of  metal  are  joined 
by  covering  their  ends  with  molten  solder,  and  when  the  latter  is 
in  a  plastic  condition  it  is  worked  into  shape  with  a  hot  tool,  which 
is  usually  made  of  copper;  a  flux  (521)  is  ordinarily  used  to  keep 
the  surfaces  to  be  joined  free  from  oxides  so  that  they  can  alloy 
with  the  solder.  The  tensile  strength  of  solder  is  greatest  when  it 
contains  72.5  per  cent  lead,  but  an  alloy  of  this  composition  has 


a 


CO  W         o 

28"    1 

GO       J> 
•+J 

"8 


fii  1! 
I  a] 

H      fl  5 


I 


PHYSICAL  PROPERTIES  OF  METALS.    ALLOYS          453 

such  a  high  freezing-point  that  it  cannot  be  readily  worked  with  a 
soldering  tool.  Alloys  of  lead  and  tin  containing  as  high  as  70 
per  cent  of  the  former  are  used  for  coating  sheet  iron  or  steel  for 
roofing,  for  castings,  etc.  Plumber's  solder  contains  about  67 
per  cent  of  lead  and  33  per  cent  of  tin;  it  begins  to  assume  the 
plastic  state  at  about  245°  and  remains  in  this  condition  during  a 
fall  in  temperature  of  65  degrees.  While  in  this  condition  the  joint 
being  made  with  the  solder,  between  two  pipes  for  example,  is 
" wiped" — that  is,  the  plastic  mass  is  molded  into  the  desired  form. 

540.  Ordinary  soft  or   tinner's   solder  contains   approximately 
50  per  cent  of  each  metal.     It  can  be  worked  at  a  low  temperature 
(see  diagram)  and  is  used  for  making  joints  which  do  not  have  to  be 
"  wiped,"  such  as  those  in  tin  cans.     It  is  more  expensive  than 
plumber's  solder  because  the  price  of  tin  is  much  higher  than  that  of 
lead.    In  soldering  tin,  an  alloy  is  used  consisting  of  a  mixture  of 
2  parts  tin,  2  parts  lead,  and  1  part  bismuth;   the  latter  is  added 
because  it  lowers  the  melting-point  of  the  solder  to  140°  so  that 
it  can  be  used  without  risk  of  melting  the  tin.     Plumber's  solder 
begins  to  solidify  about  40  degrees  above  the  melting-point  of  tin 
and  tinner's  solder  at  only  10  degrees  below  this  point. 

541.  The  alloys  of  tin  and  lead  have  been  discussed  at  some 
length  because  they  are  the  simplest  examples  of  alloys  of  this  type 
that  are  used  commercially.     When  three  or  more  substances  are 
used  in  making  an  alloy  the  conditions  are  much  more  complicated 
than  those  which  have  been  described,  and  a  wider  range  in  proper- 
ties can  be  obtained.    In  the  case  of  certain  metals,  compounds  of 
definite  composition  are  formed  as  the  result  of  the  union  of  the 
metals,  and  these  act  as  distinct  substances  and  by  increasing  the 
number  of  phases  present  add  to  the  complexity  of  the  resulting 
alloy.     By  selecting  the  right  metals  and  Using  them  in  the  correct 
proportions,  alloys  can  be  made  with  any  desired  degree  of  hard- 
ness,  ductility,   toughness,   tenacity,   fusibility,   etc.     For  many 
uses  they  are  to  be  preferred  to  the  pure  metals.       For  example, 
sound  castings  can  be  made  from  copper  with  difficulty,  and  on 
account  of  the  softness  of  the  metal  it  cannot  be  filed  and  machined 
easily.    Brass,  on  the  other  hand,  which  is  an  alloy  of  copper  and 
zinc,  can  be  molded  and  worked  with  facility. 

Small  amounts  of  non-metallic  elements  are  frequently  added 
to  metals  that  are  to  be  used  in  making  castings.     When  gases 


454  INORGANIC  CHEMISTRY  FOR  COLLEGES 

separate  from  molten  metals  on  solidification  they  cause  objection- 
able flaws  in  the  casting.  The  separation  of  gases  may  be  due 
to  the  fact  that  they  were  dissolved  by  the  molten  metal  and  lib- 
erated when  it  solidified,  or  that  they  were  present  in  the  impure 
metal  as  substances  that  interacted  at  the  high  temperature  to  form 
a  gas.  Commercial  copper,  for  example,  contains  traces  of  copper 
sulphide  and  copper  oxide,  which  interact  at  high  temperatures 
to  form  copper  and  sulphur  dioxide;  when  it  solidifies  the  gas  pro- 
duced in  this  way  is  expelled.  If  certain  non-metals  are  added 
to  the  molten  metal,  they  unite  with  the  oxygen  present  and  form 
oxides  that  are  not  decomposed  at  the  temperatures  used;  phos- 
phorus and  silicon  are  used  for  this  purpose.  Boron  has  recently 
been  found  to  be  of  particular  value  in  making  sound  castings 
from  copper  (518).  An  excess  of  the  non-metallic  element  above 
that  required  for  deoxidation  generally  results  in  increasing  the 
hardness  and  brittleness  of  the  alloys  in  which  it  is  used.  Active 
metals  whose  oxides  are  stable  at  high  temperatures  are  also 
used  as  deoxidizers;  among  these  are  aluminium,  magnesium,  and 
manganese. 

The  crystalline  structure  of  alloys  makes  them  susceptible  to 
mechanical  working  and  heat  treatment;  the  changes  effected  by 
these  agencies  are  more  marked  than  those  referred  to  in  the  case 
of  metals  (533). 

542.  The  peculiar  physical  structure  of  alloys  is  utilized  in 
making  so-called  anti-friction  alloys  which  are  used  as  bearings  in 
machinery.  These  alloys  consist  of  crystals  of  a  hard  constituent 
imbedded  in  a  soft  matrix;  as  a  result,  when  they  are  used  in  the 
bearings  for  a  steel  shaft  the  hard  material  readjusts  itself  in  the 
matrix  to  any  inequalities  in  the  shaft,  and  friction  is  thereby 
reduced;  the  soft  material  is  worn  away  at  the  surface  and  the 
hard  grains  bear  the  pressure.  Lead-antimony  alloys  are  used 
for  this  purpose.  The  eutectic  contains  13  per  cent  of  antimony; 
alloys  containing  a  larger  amount  of  the  latter  consist  of  crystals 
of  antimony,  which  is  hard,  embedded  in  the  softer  eutectic.  Bab- 
bitt metal,  which  is  a  tin-antimony-copper  alloy,  is  more  expen- 
sive, but  is  to  be  preferred  to  the  simpler  alloy.  The  composition 
of  a  few  alloys  of  technical  importance  is  given  in  the  table  on 
page  455.  The  figures  given  in  the  table  refer  to  the  several 
alloys  commonly  used.  Since  there  are  changes  in  properties 


PHYSICAL  PROPERTIES  OF  METALS.    ALLOYS 


455 


with  change  in  the  composition  of  alloys,  a  great  variety  of  sub- 
stances are  used  under  the  same  name.  For  example,  a  large 
number  of  different  brasses  are  made;  type-metal  varies  widely  in 
composition,  etc. 

COMPOSITION  OF  SOME  COMMON  ALLOYS 


Alloy 

Tin 

Copper 

Zinc 

Lead 

Anti- 
mony 

Misc. 

Babbit  metal 

889 

3  7 

7.4 

Bell  metal           

22 

78 

Brass 

95-60 

5-40 

Brittania  metal        

45.5 

1.5 

40 

13 

Bronze  for  bearings 

5 

64 

30 

Ni,  1 

Bronze,  for  castings  

10 

90 

German  silver              .    . 

60 

25 

Ni,  15 

Gold  coins,  U.  S  

10 

Au,  90 

Gun  metal            ... 

9 

91 

M^onel  metal 

28 

Ni,  69 

M^osaic  gold 

65 

35 

Fe,    3 

Nickel  Coins,  U.  S  

75 

Ni,  25 

Pewter 

80 

20 

Rose's  metal  

22.9 

27.1 

Bi,  50 

Silver  Coins,  U.  S  

10 

Ag,  90 

Silver  sterling 

2.5 

Ag,  97  5 

Solder,  plumber's  

33 

67 

Solder  soft               

50 

50 

Stereotype  metal 

10 

2 

70 

18 

Type  metal     

20 

60 

20 

Wood's  metal 

14 

24 

Bi  50 

Cd,  12 

EXERCISES 

1.  Name  several  uses  of  a  number  of  metals  based  on  the  fact  that  they 
are   (a)  hard,    (6)  soft,  and   (c)  have  a  high  heat  conductivity. 

2.  (a)  Calculate  the  weight  per  meter  of  a  copper  wire  and  an  aluminium 
wire  each  1  mm.  in  diameter.     (6)  Calculate  the  electrical  resistance  of  1 
meter  of  each  wire,     (c)  Calculate  the  diameter  of  a  wire  of  aluminium  1 
meter  of  which  has  the  same  resistance  as  1  meter  of  a  wire  of  copper  1  mm. 
in  diameter,     (d)  Calculate  the  weight  of  the  two  wires  described  in  c.     (e) 
To  what  extent  would  the  cost  of  aluminium  and  of  copper  be  a  factor  in 
deciding  which  metal  to  use  in  the  transmission  of  electricity? 


456  INORGANIC  CHEMISTRY  FOR  COLLEGES 

3.  Which  of  the  following  elements  would  be  good  and  which  poor  con- 
ductors of  electricity:    (a)  sodium,    (6)  calcium,    (c)  tellurium,    (d)  radium, 
(e)  iodine. 

4.  (a)  Why  is  it  more  agreeable  to  drink  a  hot  liquid  from  a  cup  made 
of  porcelain  than  from  one  made  of  aluminium?     (6)  Would  you  expect  a 
liquid  to  cool  more  rapidly  in  a  spoon   made  of  tin  or  of  silver?     Why? 
(c)  What  metal  would  be  best  for  making  lightning  rods? 

5.  (a)  What  properties  of  a  metal  determine  whether  it  could  be  used 
as  the  tip  on  a  soldering  iron?     (6)  Is  the  specific  heat  of  the  metal  a  factor? 
The  value  of  the  specific  heat  of  copper  is  0.091  and  of  iron  0.119.     From 
this  point  of  view  which  metal  is  preferable?     (c)  WThy  is  the  tip  of  a  solder- 
ing iron  relatively  large?     (d)  Why  is  a  steel  soldering  iron  used  with  a  solder 
containing  a  high  percentage  of  tin?     (e)  What  determines  what  kind  of 
solder  should  be  used  in  soldering  galvanized  iron?     A  solder  high  in  tin  is 
used  in  this  case.     What  conclusion  can  be  drawn  from  this  fact? 

6.  State  why  the  wiped  joint  of  solder  on  a  lead  pipe  has  a  frosted  appear- 
ance and  the  joint  on  a  tin  can  has  not. 

7.  To  what  class  of  alloys  would  you  expect  to  belong  the  following  pairs 
of  metals:   (a)  silver  and  gold,    (6)  lead  and  zinc,    (c)  copper  and  antimony? 


CHAPTER  XXXII 
THE  CHEMICAL  PROPERTIES  OF  METALS.     METALLURGY 

543.  The  metals  resemble  one  another  in  their  general  chemical 
behavior  with  other  substances,  but  they  differ  markedly  in  activity. 
They  also  show  differences  in  valence,  and  since  the  valence  of  a 
metal  is  an  important  factor  in  determining  the  chemical  properties 
of  the  compounds  containing  the  metal,  the  comparative  study 
of  these  elements  must  be  made  from  this  point  of  view. 

We  have  already  seen,  in  a  general  way,  that  the  relative  activ- 
ity of  metals  is  indicated  by  their  position  in  the  electromotive 
series  (252) ,  and  that  their  valence  follows  from  their  position  •  in 
the  periodic  classification  of  the  elements.  These  two  generaliza- 
tions will  be  of  the  greatest  value  in  correlating  the  many  facts 
in  regard  to  the  chemistry  of  the  metals  and  their  compounds, 
and  they  will  be  constantly  used  for  this  purpose. 

544.  Behavior    of    Metals    with    Other    Elements.— All    the 
metals  unite  with  the  halogens  to  form  salts ;  all  form  oxides  when 
heated  in  the  air  or  oxygen  except  those  which  fall  in  the  electro- 
motive series  below  mercury;   oxides  of  these  inactive  metals  are 
known,  however,  but  they  must  be  prepared  from  salts.     Sulphur 
unites  with  all  the  metals  to  form  sulphides;  at  high  temperatures 
most  of  the  metals  react  with  carbon  and  silicon. 

545.  Behavior  of  the  Metals  with  Water. — In  their  relative 
activity  with  water,  the  metals  fall  approximately  into  the  same 
order  as  they  do  in  the  electromotive  series.     The  alkali  metals, 
which  constitute  the  first  family  in  the  first  group  in  the  periodic 
classification,  react  rapidly  with  water  at  ordinary  temperatures, 
and  hydrogen  and  the  hydroxides  of  the  metals  are  formed.     The 
metals  of  the  alkaline  earths — those  of  family  1,  group  2 — also 
decompose  water  at  room-temperature,  but  less  rapidly. 

Magnesium  and  aluminium  react  with  cold  water  very  slowly,  but 
both  metals  can  be  activated  by  treating  them  with  a  solution  of  a 

457 


458  INORGANIC  CHEMISTRY  FOR  COLLEGES 

mercury  salt;  the  more  active  metals  set  free  mercury  and  when 
they  become  coated  with  a  trace  of  the  latter  they  decompose 
water;  hydrogen  is  evolved  and  the  metals  are  converted,  into 
hydroxides.  Such  a  combination  of  metals  is  called  a  metallic 
couple.  Zinc,  which  is  farther  down  in  the  electromotive  series, 
does  not  attack  water  at  ordinary  temperatures,  but  when  it  is 
coupled  with  copper,  by  treating  the  metal  with  the  dilute  solution 
of  a  copper  salt,  it  reacts  slowly  with  water. 

Lead,  which  is  in  the  electromotive  series  just  above  hydrogen, 
passes  into  solution  to  a  very  slight  degree  as  lead  hydroxide, 
but  as  soon  as  a  trace  of  the  latter  is  present  the  reaction  stops. 
The  metals  below  hydrogen  in  the  electromotive  series  do  not 
decompose  water. 

The  rate  at  which  metals  react  with  water  is  determined  by 
their  activity  and  by  the  temperature;  the  metals  of  the  alkalies 
and  alkaline  earths  react  more  or  less  rapidly  with  water  at  room- 
temperature;  magnesium  decomposes  water  slowly  at  100°,  zinc 
at  a  higher  temperature,  and  iron  at  red  heat. 

646.  Behavior  of  the  Metals  with  Acids. — The  behavior  of  the 
metals  when  brought  into  contact  with  acids  varies  widely;  it  is  a 
subject  of  great  importance  on  account  of  the  fact  that  it  deter- 
mines in  many  cases  what  metals  can  and  what  ones  cannot  be  used 
for  a  specific  purpose.  Ordinary  air  contains  carbon  dioxide  and 
water,  which  act  as  an  acid,  and  wherever  coal  is  burned  sulphurous 
acid  and  sulphuric  acid  are  found  in  appreciable  quantities  in  the 
air.  These  acids  attack  many  metals  and  their  presence  leads  to 
the  corrosion  of  such  metals  when  they  are  exposed  to  the  air. 

A  number  of  facts  in  regard  to  the  action  of  acids  on  metals  have 
already  been  given ;  these  will  be  restated  here  in  connection  with 
a  more  systematic  discussion  of  the  subject.  The  action  of  acids 
on  metals  is  determined  by  a  number  of  factors,  the  most  important 
of  which  are  the  following:  the  chemical  activity  of  the  metal,  the 
oxidizing  power  of  the  acid,  the  presence  of  water  which  deter- 
mines whether  or  not  the  reaction  is  one  involving  ions,  the  con- 
centration of  the  solution  of  the  acid,  the  chemical  properties  and 
solubilities  of  the  resulting  salt,  and  the  temperature. 

547.  Action  of  Non-oxidizing  Acids.  Aqueous  solutions  of 
non-oxidizing  acids  react  with  the  metals  in  the  electromotive 
series  down  to  hydrogen.  Under  these  conditions  the  reactions 


THE  CHEMICAL  PROPERTIES  OF  METALS.     METALLURGY     459 

are  ionic,  and  hydrogen  and  a  salt  of  the  metal  are  formed. 
Changes  in  concentration  and  temperature  have  the  effect  that 
might  be  expected  from  a  knowledge  of  the  influence  of  these 
factors  on  the  ionization  of  the  acid  used.  Oxidizing  acids,  if 
sufficiently  dilute,  behave  in  the  same  way,  because  their  oxidizing 
power  decreases  with  increased  dilution  as  the  result  of  the  con- 
version of  molecules — the  active  oxidizing  agent — into  ions. 
Many  metals  have  two  valencies;  iron,  for  example,  forms  two 
chlorides,  ferrous  chloride,  FeCk,  and  ferric  chloride,  FeCla,  in 
which  the  valence  of  the  element  is  2  and  3,  respectively.  Since 
iron  in  the  second  compound  is  in  a  higher  state  of  oxidation,  and 
since  nascent  hydrogen  is  an  active  reducing  agent,  we  would  not 
expect  ferric  chloride  to  be  formed  in  the  presence  of  the  hydrogen 
set  free  when  iron  dissolves  in  hydrochloric  acid;  and  this  is 
the  case.  In  general,  when  a  metal  forms  two  classes  of  salts  in 
which  it  shows  different  valencies,  the  salt  containing  the  metal  in 
the  lower  valence  is  formed  when  it  dissolves  in  dilute  acids. 

Whether  or  not  a  metal  dissolves  in  an  acid  is  determined  not 
only  by  its  activity,  but  by  the  behavior  with  water  of  the  product 
formed.  Aluminium,  a  very  active  metal,  which  stands  high  in 
the  electromotive  series,  is  scarcely  attacked  by  dilute  sulphuric, 
nitric,  and  most  other  acids.  We  shall  see  later  that  aluminium 
is  a  trivalent  element,  and  on  account  of  this  fact  its  hydroxide 
is  a  very  weak  base;  its  salts  are,  therefore,  hydrolyzed  in  water, 
and  when  the  metal  is  treated  with  a  dilute  solution  of  an  acid  the 
salt  first  formed  is  probably  decomposed  by  the  water  present  to 
form  an  insoluble  basic  salt,  which  protects  the  surface  of  the  metal 
from  further  action.  Lead  is  slightly  attacked  by  cold  dilute  solu- 
tions of  hydrochloric  acid  or  sulphuric  acid,  because  the  chloride 
and  sulphate  of  the  metal  are  difficultly  soluble  in  water,  and  the 
small  amount  first  formed  protects  the  metal  from  further  action. 
Lead  chloride  is  soluble  in  hot  water,  and  for  this  reason  the  metal 
dissolves  in  a  hot  solution  of  hydrochloric  acid;  lead  sulphate  is 
practically  insoluble  in  hot  water  and  the  metal  is  not  attacked 
by  a  hot  dilute  solution  of  sulphuric  acid.  Since  the  chlorides  of 
all  the  metals  in  the  electromotive  series  down  to  lead  are  sol- 
uble in  water,  all  these  metals  dissolve  readily  in  hydrochloric  acid. 
They  all  dissolve  in  dilute  sulphuric  acid  except  aluminium,  and 
in  dilute  nitric  acid,  except  aluminium  and  chromium. 


460  INORGANIC  CHEMISTRY  FOR  COLLEGES 

648.  Action  of  Oxidizing  Acids. — Whether  or  not  a  metal  dis- 
solves in  an  oxidizing  acid  depends  on  the  activity  of  the  metal 
and  of  the  oxidizing  acid,  the  concentration  'of  the  latter,  and  the 
properties  of  the  products  formed.  Since  oxidations  are  essen- 
tially molecular  reactions,  the  smaller  the  concentration  of  water 
present,  which  ionizes  the  acid,  the  greater  is  the  activity  of  the 
oxidizing  agent.  All  the  metals  in  the  electromotive  series  down 
to  and  including  silver  are  oxidized  by  concentrated  nitric  acid 
and  by  concentrated  sulphuric  acid.  In  certain  cases  the  results 
obtained  with  nitric  acid  appear  to  be  anomalous.  When  iron 
is  placed  in  fuming  nitric  acid  it  does  not  dissolve,  but  is  changed 
in  such  a  way  that  it  does  not  exhibit  the  properties  of  the  element 
in  its  normal  condition.  After  treatment  with  the  acid  iron  will 
not  dissolve  in  dilute  acids,  or  precipitate  copper  from  solutions 
of  its  salts;  it  is  said  to  be  passive.  If  the  metal  is  scratched  or 
struck,  it  assumes  its  usual  condition.  Chromium  and  cobalt  can 
be  rendered  passive  in  the  same  way.  It  is  probable  that  these 
metals  are  converted  on  the  surface  into  oxides  which  contain  a 
large  percentage  of  oxygen  and  do  not  show  basic  properties  and, 
therefore,  are  insoluble  in  acids;  as  they  cover  the  metal  they 
prevent  it  from  interacting  with  solutions  of  salts.  Much  work 
has  been  devoted  to  the  study  of  passive  metals  and  a  final 
conclusion  as  to  the  cause  of  the  phenomenon  has  not  been  reached. 
Aluminium  is  practically  unattacked  by  concentrated  nitric  acid. 

In  dilute  solutions  nitric  acid  acts  towards  metals  largely  in 
its  capacity  as  an  acid;  hydrogen  and  the  nitrate  containing  the 
metal  in  its  lower  valence  are  formed;  iron  and  dilute  nitric  acid 
yield  ferrous  nitrate,  Fe(NOs)2,  and  hydrogen.  With  increasing 
concentration  the  oxidizing  property  of  nitric  acid  shows  itself; 
with  strong  nitric  acid  iron  forms  ferric  nitrate,  Fe(NO3)3,  and 
as  the  result  of  its  acting  as  an  oxidizing  agent  the  acid  is  reduced 
to  nitric  oxide  or  other  oxides  of  nitrogen.  Concentrated  nitric 
acid  has  little  effect  on  cast  iron,  and  its  vapor  less;  as  a  conse- 
quence, retorts  made  of  cast  iron  are  used  in  manufacturing  the 
acid  from  sodium  nitrate,  and  care  is  taken  to  have  the  tempera- 
ture of  the  retort  above  the  boiling-point  of  the  acid.  Since 
there  is  a  slight  action  between  the  acid  and  the  metal,  the  con- 
densers used  are  made  of  aluminium,  glass,  or  other  resisting  non- 
metallic  material. 


THE  CHEMICAL  PROPERTIES  OF  METALS.    METALLURGY     461 

The  action  of  nitric  acid  on  tin  illustrates  another  point.  The 
metal  forms  compounds  in  which  it  functions  in  its  metallic  char- 
acter and  shows  the  valence  2;  it  also  forms  compounds  in  which 
the  valence  is  4,  and  in  which  it  plays  the  part  of  an  acid-forming 
element.  Stannic  oxide,  SnO2,  is  an  acid  anhydride  and  yields 
salts  known  as  stannates.  Tin  dissolves  in  dilute  nitric  acid  and 
stannous  nitrate  is  formed ;  when  treated  with  concentrated  nitric 
acid  the  metal  is  oxidized  to  the  dioxide,  which  does  not  dissolve 
in  the  acid,  and  as  a  result  a  white  insoluble  powder,  SnO2,£H2O 
is  formed. 

Since  the  normal  nitrates  are  all  soluble  and  those  of  the  weakly 
metallic  elements  are  not  hydrolyzed  in  the  presence  of  a  large 
excess  of  nitric  acid,  all  the  metals  except  those  noted  above  dis- 
solve in  the  strong  acid. 

Hot  concentrated  sulphuric  acid  oxidizes  all  the  metals  down 
to  and  including  silver.  The  solubility  of  the  sulphate  formed  has 
a  marked  effect  on  the  extent  to  which  the  reaction  takes  place. 
Lead  sulphate,  for  example,  is  practically  insoluble  in  sulphuric 
acid  containing  water  until  the  concentration  of  the  acid  reaches 
about  77  per  cent;  it  is,  accordingly,  not  appreciably  attacked  by 
acid  up  to  this  concentration  because  the  sulphate  first  formed 
serves  as  a  coating  to  protect  the  metal;  it  dissolves  when  heated 
with  stronger  acid  because  the  sulphate  formed  passes  into  solu- 
tion. The  use  of  lead  in  the  construction  of  chemical  apparatus 
in  which  sulphuric  acid  is  to  be  used  is  limited  by  these  facts. 

The  behavior  of  iron  with  sulphuric  acid  is  quite  different  from 
that  of  lead;  the  sulphate  of  the  metal  is  soluble  in  water  and, 
consequently,  the  metal  is  attacked  by  mixtures  of  the  acid  and 
water  until  the  concentration  of  the  former  reaches  about  77  per 
cent,  when  action  practically  ceases.  Application  is  made  of  this 
fact  in  the  concentration  of  the  sulphuric  acid  made  in  the  cham- 
ber process;  the  acid  is  concentrated  in  lead  pans  until  it  reaches 
77  per  cent  and  is  then  transferred  to  iron  stills  in  which  the  con- 
centration to  98  per  cent  is  continued;  it  has  been  found  that 
cast  iron  serves  best. 

649.  The  Effect  of  the  Presence  of  Oxygen  on  the  Solubility  of 
Metals  in  Acids. — In  the  presence  of  air  solutions  of  acids  dis- 
solve certain  metals  that  are  not  attacked  by  these  acids  in  the 
absence  of  oxygen.  It  is  probable  that  solution  results  from  the 


462  INORGANIC  CHEMISTRY  FOR  COLLEGES 

fact  that  the  metal  is  first  oxidized  by  the  oxygen  and  the  oxide 
formed  dissolves  in  the  acid  present.  Copper,  for  example,  which 
is  below  hydrogen  in  the  electromotive  series,  does  not  react  with 
dilute  acids,  but  if  oxygen  is  present  it  is  more  or  less  rapidly 
converted  into  a  salt.  Advantage  of  this  fact  is  taken  in  one  of 
the  commercial  methods  for  the  preparation  of  copper  sulphate. 
Scrap  copper  is  allowed  to  stand  with  dilute  sulphuric  acid  in 
contact  with  the  air  until  it  passes  into  solution.  Although  the 
reaction  takes  place  slowly,  it  is  more  economical  and  convenient 
than  dissolving  the  metal  in  hot  concentrated  sulphuric  acid  with 
the  attendant  loss  of  sulphur  dioxide,  the  difficulty  of  handling 
the  hot  acid,  and  the  necessity  of  using  a  non-metallic  container 
for  the  latter. 

Copper  reacts  with  dilute  acetic  acid  in  the  presence  of  air; 
the  basic  salt  formed,  which  is  insoluble  in  water,  is  used  as  a 
green  pigment  under  the  name  verdigris.  Tin,  which  is  low  in  the 
electromotive  series,  near  hydrogen,  is  scarcely  attacked  by  dilute 
solutions  of  weak  acids;  in  the  presence  of  oxygen,  however,  it  is 
dissolved. 

550.  The  Corrosion  of  Metals  in  the  Presence  of  Air. — The 
metals  down  to  and  including  copper  are  more  or  less  affected  by 
contact  with  the  air;  either  an  oxide  is  formed  or,  in  certain  cases, 
the  presence  in  the  air  of  carbon  dioxide  leads  to  the  production 
of  insoluble  basic  carbonates.  Whether  or  not  the  metal  is  appre- 
ciably corroded  is  determined  by  the  physical  properties  of  the 
rust  formed.  In  the  case  of  aluminium,  for  example,  the  super- 
ficial coating  of  hydrated  oxide  adheres  firmly,  and  as  it  is  very 
thin  it  scarcely  alters  the  appearance  of  the  metal;  with  mag- 
nesium, however,  another  active  metal,  the  result  is  different,  for 
the  basic  carbonate  formed  is  light  and  does  not  adhere  so  firmly, 
and,  as  a  consequence,  the  metal  is  slowly  disintegrated.  Nickel, 
cobalt,  and  tin  are  affected  to  only  a  slight  degree  if  at  all;  zinc 
becomes  covered  with  a  thin  coating  of  basic  carbonate,  which 
resists  further  action  and  is,  therefore,  permanent  in  the  air. 
Lead  becomes  coated  with  a  very  thin  adhering  layer  of  basic 
carbonate  and  resists  further  corrosion.  A  clean  surface  of  copper 
oxidizes  rapidly  in  the  air,  but  the  coating  of  the  oxide  formed 
is  so  thin  it  only  slightly  deepens  the  color  of  the  metal;  it  is  very 
permanent  in  ordinary  air. 


THE  CHEMICAL  PROPERTIES  OF  METALS.    METALLURGY     463 

The  behavior  of  iron  is  quite  different;  in  this  case  the  oxide 
formed  is  a  powder  the  density  of  which  is  less  than  that  of  iron; 
when  it  is  formed  it  occupies  more  space  than  the  metal  from 
which  it  was  produced,  and,  therefore,  falls  away  from  the  sur- 
face, leaving  freshly  exposed  metal  to  be  acted  on  farther;  and 
the  rusting  continues  more  or  less  rapidly. 

A  great  deal  of  attention  has  been  devoted  to  the  study  of  the 
corrosion  of  iron  on  account  of  its  technical  importance;  some  of 
the  more  important  conclusions  reached  will  be  discussed  later 
(755). 

In  the  presence  of  the  air  of  cities,  especially  near  places  where 
large  quantities  of  coal  are  burned,  the  corrosion  of  metals  takes 
place  much  more  rapidly  than  in  the  pure  air  of  the  country.  The 
sulphuric  acid  in  the  air,  which  is  produced  from  the  sulphur  in  coal, 
attacks  the  metals  in  the  presence  of  oxygen,  and  salts  are  formed 
which  are  eventually  converted  in  the  presence  of  water  and  carbon 
dioxide  into  basic  carbonates.  The  green  color  of  bronze  statuary 
kept  out  of  doors  and  of  copper  roofing  is  due  to  this  cause.  In 
the  neighborhood  of  the  sea,  copper  and  even  nickel,  which  resists 
ordinary  air,  become  covered  with  a  coating  of  green  basic  car- 
bonates. The  corrosion  under  these  circumstances  is  due  to 
the  fact  that  the  ocean  and  the  air  in  its  neighborhood  contain 
salts  that  hydrolyze  slightly  and  furnish  a  trace  of  hydrochloric 
acid  which  attacks  the  metal. 

Hydrogen  sulphide  is  present  in  the  air  under  certain  condi- 
tions, and  as  it  is  decomposed  by  most  metals  with  the  formation 
of  hydrogen  and  a  metallic  sulphide,  its  presence  leads  to  the  tar- 
nishing of  bright  metallic  surfaces.  Copper  and  silver  act  rapidly 
with  the  gas,  and  as  the  sulphides  of  these  metals  are  black  they 
are  quickly  tarnished.  When  the  layer  of  sulphide  formed  on 
silver  is  very  thin  it  gives  to  the  metal  a  yellow  color. 

551.  The  facts  enumerated  above  in  regard  to  the  corrosion 
of  metals  are  taken  into  account  in  the  use  of  metals  for  indus- 
trial purposes.  Copper  and  zinc  are  used  in  the  construction  of 
buildings  when  it  is  necessary  to  expose  a  metal  to  the  air.  Roofs 
were  formerly  covered  with  copper  or  lead,  but  the  high  price  of 
the  metals  has  resulted  in  the  substitution  of  cheaper  material  for 
this  purpose.  The  question  of  cost  has  been  met  in  many  cases  by 
covering  iron  or  steel  with  a  coating  of  a  resistant  metal.  Gal- 


464  INORGANIC  CHEMISTRY  FOR  COLLEGES 

vanized  iron  is  made  by  dipping  iron  into  molten  zinc ;  the  metal 
that  adheres  firmly  to  the  iron  if  it  has  been  carefully  cleaned, 
serves  to  protect  the  latter  for  a  long  time  from  corrosion.  In 
the  process  known  as  "  sherardizing  "  objects  made  of  iron,  after 
treatment  with  dilute  sulphuric  acid  to  remove  all  oxide,  are  baked 
in  zinc  dust,  and,  as  a  result,  are  covered  with  a  layer  of  the  metal. 
Tin-plate  is  prepared  by  dipping  sheet-iron  into  molten  tin;  iron 
is  also  coated  with  lead  in  a  similar  way. 

552.  Nickel,  silver,  zinc,  gold,  and  copper  are  applied  to  the 
surfaces  of  other  metals  by  the  process  of  electroplating,  which 
will  be  described  later.     Other  substances  than  metals  are  used 
as  protective  coatings  to  prevent  corrosion.     Transparent  lacquers 
and  varnishes  made  from  organic  gums  are  frequently  used  to 
protect  iron  and  silver.     When  iron  is  heated  to  a  high  tempera- 
ture with  steam  it  reacts  and  an  oxide  of  the  formula  Fes 04  is 
formed  (736).     This  fact  is  used  in  covering  articles  made  of  iron 
or  steel  with  the  oxide,  which  acts  as  a  protective  coating;  when 
they  are  heated  in  a  closed  retort  and  subjected  to  the  action  of 
high-pressure  steam  there  is  formed  a  surface  layer  of  the  oxide, 
which  has  a  blue  color.     The  process  is  used  with  rifle  barrels, 
pistols,  the  so-called  Russia  or  black  sheet  iron,  etc.     Paints  of 
various  kinds  are  also  used  to  protect  iron  from  corrosion.     The 
protection  against  corrosion   afforded  by  protective  coatings  is 
dependent  on  the  physical  nature  of  the  latter;  if  they  are  porous 
or  contain  minute  holes,  air  will  reach  the  metal  underneath  and 
corrosion  will  take  place  more  or  less  rapidly.     It  requires  skill 
and  attention  to  details  to  galvanize  or  tin  iron  properly  and,  as  a 
consequence,  many  commercial  articles  of  iron  protected  in  this 
way  soon  rust. 

553.  As  the  result  of  the  intensive  study  of  alloys  in  recent 
years  many  substances  of  this  class  have  been  made  that  resist 
corrosion  by  the  air  and  chemicals  in  general,  and  they  are  exten- 
sively used  to  replace  metals.     An  alloy  of  iron  and  silicon  known 
as  duriron  is  particularly  resistant  and  has  found  many  uses, 
especially  in  chemical  factories.     The  art  of  enameling  metals 
has  also  developed,  recently,  and  many  large  pieces  of  apparatus 
constructed  of  metal  are  lined  with  this  type  of  material,  which 
resists  the  action  of  solutions  of  most  chemicals.     Of  the  chemicals 
ordinarily  used,  hydrochloric  acid  is,  perhaps,  the  most  difficult 


THE  CHEMICAL  PROPERTIES  OF  METALS.     METALLURGY     465 

to  protect  against,  because  it  is  a  very  active  acid  and  forms  sol- 
uble salts  with  most  metals. 

554.  The  Action  of  Metals  on  Alkalies. — The  hydroxides  of 
the  metallic  elements  are  bases,  for  they  dissolve  in  acids  and 
form  salts;  but  those  of  certain  metals  act  also  as  acids,  dissolve 
in  bases,  and  form  salts  in  which  the  metal  plays  the  part  of  an 
acid-forming  element.     For  example,  zinc  hydroxide  can  function 
either  as  an  acid  or  a  base;   with  hydrochloric  acid  it  forms  zinc 
chloride  and  with  sodium  hydroxide,  sodium  zincate,  Zn(ONa)2 
or  Na2ZnO2.     Most  metals  of  this  class  react  with  active  bases; 
examples  of  these  are  aluminium,  zinc,  tin,  and  lead.     Arsenic 
and  antimony,  which  are  essentially  acid-forming  elements,  are 
also  soluble  in  alkalies. 

Certain  metals  which  do  not  react  with  solutions  of  bases  do 
react  with  the  alkalies  at  the  temperature  of  fusion;  platinum, 
for  example,  is  converted  slowly  under  these  conditions  into  a 
platinate,  and  gold  into  an  aurate.  Other  metals  that  resist 
fused  alkalies  are  dissolved  if  an  oxidizing  agent,  such  as  sodium 
nitrate,  is  present.  "  Owing  to  the  presence  of  the  latter  the  metal 
is  oxidized  to  its  higher  valence  in  which  it  possesses  acidic  proper- 
ties, and  it  passes  into  solution  as  a  result;  manganese,  chromium, 
and  iron,  examples  of  metals  that  behave  in  this  way,  yield  man- 
ganates,  chromates,  and  ferrates. 

Nickel  does  not  show  acidic  properties  in  any  of  its  compounds 
and  as  it  melts  at  a  relatively  high  temperature  it  is  used  in  making 
vessels  to  be  employed  in  the  preparation  of  compounds  involving 
fusion  with  an  alkali  and  an  oxidizing  agent.  If  there  is  no  chance 
for  oxidation,  iron  can  be  used  for  the  construction  of  apparatus 
in  which  fusions  with  alkalies  are  carried  out. 

555.  Action  of  Metals  on  Salts. — The  replacement  of  one  ele- 
ment by  another  when  a  metal  is  placed  in  a  solution  of  a  salt  has 
already  been  discussed  (252).     It  will  be  recalled  that  whether 
or  not  a  reaction  takes  place  is  determined  by  the  relative  tenden- 
cies of  the  two  metals  to  form  ions.     In  general,  when  there 
are  no  disturbing  factors  present,  a  metal  will  displace  a  second 
metal  from  its  salts  if  the  first  metal  is  above  the  second  in  the 
electromotive  series.     A  number  of  applications  are  made  of  this 
property.     For  example,  pins,  which  are  made  of  brass  or  iron,  are 
coated  with  tin  by  placing  them  in  a  solution  of  a  tin  salt.     The 


466  INORGANIC  CHEMISTRY  FOR  COLLEGES 

action  of  metals  on  fused  salts  is  utilized  in  isolating  certain  metals 
that  cannot  be  readily  separated  from  their  oxides  by  reduction 
with  carbon.  Aluminium  was  first  prepared  in  this  way  by  heating 
sodium  with  fused  aluminium  chloride.  When  salts  are  in  the 
molten  state  they  show  the  electrical  properties  that  they  exhibit 
in  aqueous  solutions;  they  conduct  electricity  and  enter  into 
reactions  similar  to  those  which  take  place  as  the  result  of  the 
presence  of  ions.  As  a  consequence,  the  replacement  of  metals, 
one  by  the  other,  when  metals  are  brought  into  contact  with  fused 
salts,  resembles  closely  the  replacement  when  solutions  of  salts 
are  used. 

556.  Occurrence  of  the  Metals.— F.  W.  Clarke  of  the  U.  S. 
Geological  Survey  has  calculated  from  all  available  data  the 
probable  composition  of  the  earth's  crust  and  atmosphere  as  far 
as  it  has  been  possible  to  examine  it  on  its  surface,  at  great  heights 
in  the  air,  and  at  the  depth  obtained  in  deep  mines.  The  results 
are  given  in  the  following  table: 

COMPOSITION  OF  THE  EARTH'S  CRUST 

Per  Cent  Per  Cent 

Oxygen 49 . 85      Titanium 0.41 

Silicon 26 . 03      Chlorine 0 . 21 

Aluminium 7 . 28      Carbon 0 . 19 

Iron 4. 12      Phosphorus 0. 10 

Calcium 3. 18      Fluorine 0. 10 

Sodium 2.33      Sulphur 0.10 

Potassium 2.33      Barium 0.09 

Magnesium 2. 11      Manganese 0.08 

Hydrogen 0.97      Nitrogen 0.03 


.       99.51 

From  the  above  it  is  seen  that  the  total  of  all  the  elements  other 
than  those  listed  is  less  than  one-half  per  cent.  Aluminium  is 
widely  distributed  and  occurs  chiefly  as  aluminium  silicate  in 
association  with  other  silicates,  which  make  up  the  more  important 
rocks.  As  the  result  of  the  disintegration  of  the  latter  under 
atmospheric  influences,  the  soil  is  formed.  As  a  consequence,  it  is 
clear  why  oxygen,  silicon,  and  aluminium  occupy  their  positions 
in  the  above  table. 

The  other  metals  listed  in  the  table  occur  as  silicates  in  igneous 


THE  CHEMICAL  PROPERTIES  OF  METALS.    METALLURGY     467 

rocks,  and  also  in  other  forms,  which  owe  their  origin  and  existence 
to  the  chemical  properties  of  the  metals,  and  to  the  physical  prop- 
erties of  the  salts  formed  from  them.  When  the  earth  was  being 
formed  as  the  result  of  the  cooling  of  a  mass  made  up  of  elements  in 
the  form  of  vapor,  the  more  active  elements  probably  united  first 
and  the  chlorides  of  the  alkali  metals  and  alkaline  earths  were 
soon  formed.  Since  these  compounds  are  soluble  in  water  they 
have  accumulated  in  the  ocean  and  in  salt  beds  formed  as  the 
result  of  the  drying  up  of  inland  seas.  Other  soluble  salts  of  the 
alkali  metals  occur  in  nature  in  arid  regions;  the  sodium  nitrate 
of  Chile  is,  in  all  probability,  the  result  of  the  nitrification  of  organic 
matter  which  was  accumulated  in  one  locality.  Sodium  car- 
bonate is  found  in  certain  alkaline  lakes,  and  borax  in  desert 
regions. 

The  alkaline  earths  occur  as  carbonates  and  sulphates,  calcium 
carbonate  being  an  important  constituent  of  many  rocks,  such 
as  limestone  and  marble;  gypsum,  CaSO4,2H2O,  is  also  an 
important  mineral.  Magnesium  occurs  extensively  as  silicates, 
as  the  sulphate  in  salt  deposits,  and  as  the  carbonate  in  rocks. 

As  we  pass  to  the  less  active  metals  in  the  electromotive  series, 
in  addition  to  silicates,  we  find  carbonates,  oxides,  hydrated 
oxides,  and  sulphides  as  the  chief  minerals  containing  these  ele- 
ments. When  we  reach  copper  in  descending  in  the  series,  we  find 
that  the  metal  occurs  in  the  above  forms  but  is  also  found  in  the 
free  condition  in  nature.  Metallic  iron,  cobalt,  and  nickel  are 
present  in  some  meteors,  but  they  do  not  occur  in  this  form  in  the 
earth's  crust.  The  elements  below  copper  in  the  series  all  occur 
in  the  native  state,  that  is,  not  in  combination  with  another 
element.  The  more  active  of  these  are  also  found  in  combina- 
tion, sulphides  being  the  most  abundant  of  these  compounds. 
Silver  chloride  is  an  important  mineral.  The  noble  metals  occur 
almost  exclusively  in  the  metallic  condition;  gold,  however, 
occurs  as  the  sulphide  and  telluride. 

METALLURGY 

557.  The  science  which  treats  of  the  methods  used  to  obtain 
the  metals  in  the  free  condition  from  the  compounds  that  occur  in 
nature,  is  called  metallurgy.  All  the  naturally  occurring  com- 


468  INORGANIC  CHEMISTRY  FOR  COLLEGES 

pounds  cannot  be  used  as  sources  of  the  metals  on  account  of  the 
difficulties  encountered  in  certain  cases.  Potassium,  for  example, 
is  widely  distributed  in  the  form  of  feldspar  and  other  silicates, 
but  these  cannot  be  used  economically  as  a  source  of  the  metal  or 
its  compounds;  potassium  chloride,  however,  can  be  readily  used 
for  these  purposes.  The  minerals,  in  more  or  less  pure  form, 
which  are  actually  used  as  sources  of  the  metals,  are  known  as 
ores.  These  are  commonly  oxides,  sulphides,  carbonates,  and 
chlorides — compounds  simple  in  composition,  from  which  the 
metal  can  be  extracted  by  the  use  of  a  few  chemical  reactions 
comparatively  easy  to  carry  out. 

The  process  employed  in  the  isolation  of  any  particular  metal 
is  determined  by  the  chemical  properties  of  the  compound  used 
and  those  of  the  impurities  present  in  the  ore.  If  we  take  iron 
as  an  example,  its  chief  ores  are  oxides  of  the  metal  mixed  with 
sand  and  other  silicious  material.  Since  the  oxides  of  iron  are 
reduced  to  the  metal  by  carbon  at  a  high  temperature,  the  ore 
is  mixed  with  coke  and  heated  in  a  blast  furnace  to  the  tem- 
perature at  which  iron  melts.  A  blast  of  air  is  blown  into  the 
furnace  during  the  reduction  to  burn  a  part  of  the  coke  and 
thus  produce  the  heat  required.  Since,  at  the  temperature  used, 
the  silicon  dioxide  present  in  the  ore  does  not  melt  and  in  the 
solid  form  prevents  contact  between  the  oxide  of  iron  and  the 
reducing  agent,  calcium  carbonate  is  added  to  the  charge  of 
ore  and  coke;  it  forms  a  silicate  with  the  silicon  dioxide,  and  all 
the  silicious  material  melts.  The  calcium  carbonate  is,  thus, 
used  as  a  flux;  the  molten  material  other  than  the  iron  is  called 
the  slag. 

A  number  of  metals  which  occur  as  oxides  are  obtained  by 
reducing  their  ores  with  carbon;  among  these  are  zinc,  iron,  tin, 
and  bismuth. 

558.  When  the  ore  to  be  used  is  a  sulphide,  it  is  converted  into 
an  oxide  before  reduction;  carbon  does  not  reduce  sulphides, 
because  the  reaction  between  carbon  and  sulphur  is  endothermic. 
The  ore  is  first  heated  in  the  air,  or  roasted,  as  a  result  of  which  it 
is  changed  into  an  oxide,  and  sulphur  dioxide  is  given  off;  the 
oxide  is  then  reduced  by  carbon.  Zinc  is  obtained  in  this  way 
from  its  sulphide,  sphalerite,  ZnS.  Other  elements  obtained  in 
this  way  from  sulphides  are  cadmium,  cobalt,  and  nickel. 


THE  CHEMICAL  PROPERTIES  OF  METALS.     METALLURGY     469 

In  the  case  of  certain  metals  whose  oxides  are  readily  reducible, 
it  is  not  necessary  to  use  carbon  as  a  reducing  agent  when  the  ore 
is  a  sulphide.  Advantage  is  taken  of  the  fact  that  the  oxide  is 
reduced  by  the  sulphur  in  combination  with  the  metal  as  sulphide. 
For  example,  when  copper  sulphide  is  heated  with  copper  oxide 
the  two  compounds  interact  and  copper  is  formed: 

CuS  +  2Cu2O  =  5Cu  4-  SO2 

In  the  case  of  metals  of  this  kind — copper  and  lead  are  examples — 
the  ore  is  roasted  until  only  the  proper  amount  of  it«is  converted 
into  oxide;  it  is  then  heated  to  a  higher  temperature  and  as  the 
result  of  the  interaction  of  the  oxide  and  sulphide  the  metal  is 
obtained  in  the  free  condition. 

559.  Silver  usually  occurs  with  lead  or  copper;   its  separation 
from  these  metals  after  extraction  from  their  ores  will  be  described 
later  (713). 

The  isolation  of  mercury  from  its  sulphides  is  easy  to  effect 
because  they  are  converted  into  the  oxide  when  roasted,  and  the 
latter  readily  breaks  down  at  a  higher  temperature  into  the  metal 
and  oxygen.  All  that  is  necessary  is  to  heat  the  ore  in  air  and 
condense  the  vapor  of  the  metal  formed. 

The  metals  of  the  platinum  group  are  found  free  in  nature  more 
or  less  alloyed  with  one  another;  they  are  separated  from  one 
another  by  converting  them  into  salts  which  differ  in  solubility. 

Gold  occurs  native,  in  the  sulphide  ores  of  lead  and  copper, 
and  as  a  telluride  along  with  silver.  Since  its  metallurgy  consists 
largely  in  its  separation  from  other  metals  it  will  be  considered 
later. 

560.  The  metals  which  have  been  discussed  up  to  this  point 
are  those  whose  oxides  are  reducible  by  carbon  at  temperatures 
readily  obtained  when  the  latter  burns.     We  have  seen  that  the 
methods  used  in  the  isolation  of  these  metals  varied  with  the 
stability  of  the  oxides,  and  that  they  changed,  in  general,  progress- 
ively with  the  position  of  the  metal  in  the  electromotive  series 
of  the  elements.     In  the  case  of  metals  whose  oxides  are  not 
reducible  by  carbon  at  the  temperature  produced  when  the  latter 
burns  and,  therefore,  obtained  in  furnaces  (about  1800°),  other 
methods  have  to  be  employed.     Either  the  reduction  is  carried 
out  in  an  electric  arc  furnace,  where  the  temperature  reaches 


470  INORGANIC  CHEMISTRY  FOR  COLLEGES 

about  3500°,  or  the  metal  is  isolated  by  the  electrolysis  of  one  of  its 
compounds  in  the  molten  condition.  In  certain  cases  a  fused  salt 
of  the  element  is  heated  with  a  more  active  element,  which  has 
been  previously  prepared  by  electrolysis. 

Calcium  can  be  obtained  by  reducing  its  oxide  with  carbon  in 
an  electric  furnace,  but  as  the  metal  unites  with  carbon  to  form 
calcium  carbide,  CaC2,  care  has  to  be  taken  to  avoid  an  excess  of 
the  reducing  agent.  The  metal  is  ordinarily  obtained  as  the 
result  of  the  electrolysis  of  fused  calcium  chloride;  magnesium  is 
isolated  in  a  similar  way  from  a  fused  double  salt,  MgCl2,KCl. 
Double  salts  are  frequently  used  in  the  electrolysis  because  they 
melt  at  a  lower  temperature  than  the  simple  compounds.  Sodium 
is  obtained  by  electrolyzing  fused  sodium  hydroxide;  and  potas- 
sium from  the  fused  chloride. 

Aluminium,  which  is  of  such  technical  importance,  is  prepared 
on  the  large  scale  by  the  electrolysis  of  a  solution  of  its  oxide  in 
fused  cryolite,  AlFs,3NaF.  Further  details  of  some  of  these 
electrolytic  processes  will  be  given  later, 

EXERCISES 

1.  Tabulate  the  properties  of  the  common  metals  given  in  Chapter  XXXII 
in  the  following  way:  Prepare  a  table  in  which  the  following  metals  arranged 
according  to  their  position  in  the  electromotive  series  form  the  first  vertical 
column:  Na,  Ca,  Mg,  Al,  Zn,  Fe,  Sn,  Pb,  Cu,  Hg,  Ag,  Au.     Place  horizontally 
the  following  headings:    (a)  heated  in  air,     (6)  air  at  room  temperature, 
(c)  water,    (d)  solutions  of  non-oxidizing  acids,    (e)  oxidizing  acids,    (/)  solu- 
tion of  alkalies,    (gr)  occurrence  in  nature,    (h)  metallurgy.     Draw  lines  hori- 
zontally between  the  syrrbols  of  the  elements  and  vertically  between  the 
headings  and  fill  in  each  square  formed  with  a  brief  statement  which  gives 
the  facts  called  for. 

2.  Making  use  of  only  the  facts  given  in  this  chapter,  how  could  you  dis- 
tinguish the  following  from  each  other:    (a)  Mg  and  Ag,    (6)  Cu  and  Au, 
(c)  SnandZn,   (d)  CaandPb,  (e)  Al  and  Zn,  (/)  Al  and  Sn,   (g)  PbandZn? 


CHAPTER  XXXIII 
ELECTROCHEMISTRY 

561.  Electrical  energy  is  used  extensively  in  the  preparation  of 
a  large  number  of  compounds  which  are  of  great  industrial  im- 
portance; and  its  use  is  being  constantly  extended  as  cheap  sources 
of  electrical  power  are  made  available  and  the  price  of  coal  rises. 
The  up-keep  and  utilization  of  an  electrical  installation  using 
water-power  is  small  compared  with  one  requiring  fuel  as  a  source 
of  energy,  for  after  the  first  cost  has  been  met,  the  energy  con- 
sumed is  derived  from  falling  water.  Electrical  energy  is  used  in 
chemical  operations  as  a  source  of  heat  in  furnaces  of  various 
types,  and  is  directly  converted  into  chemical  energy  in  the  prep- 
aration of  metals  from  their  ores,  and  in  the  manufacture  of  many 
important  compounds.  This  branch  of  industry  has  become  so 
important  and  the  knowledge  required  to  conduct  and  develop 
it  so  detailed  that  it  has  led  to  a  new  type  of  specialist  known  as 
an  electrochemical  engineer. 

The  conversion  of  electrical  energy  into  chemical  energy  can 
be  carried  out  quantitatively  in  certain  cases  in  either  direction, 
and  as  we  cannot  measure  chemical  energy  directly,  it  is  necessary 
to  use  such  transformations  in  order  to  get  an  accurate  knowledge 
of  the  energy  changes  in  chemical  reactions.  We  have  seen 
how  important  is  a  knowledge  of  these  changes  in  energy,  and  a 
chemist  must,  therefore,  be  familiar  with  those  parts  of  the  science 
of  electricity  which  are  used  in  interpreting  chemical  change. 

From  the  practical  point  of  view,  and  from  the  theoretical 
aspect  of  the  subject,  which  is  the  basis  for  applications  of  elec- 
tricity in  chemical  industry,  a  knowledge  of  electrochemistry  is  a 
necessity.  The  nature  of  the  electric  current  and  the  methods 
used  to  measure  it  will  be  first  briefly  considered,  and  then  its 
simpler  applications  in  industry  and  the  pure  science  will  be  dis- 
cussed in  some  detail. 

471 


472  INORGANIC  CHEMISTRY  FOR  COLLEGES 

562.  The  Nature  of  the  Electric  Current. — In  the  modern  con- 
ception of  the  constitution  of  matter  which  has  been  referred 
to  briefly  (398),  the  view  is  held  that  the  atom  is  made  up  of  a  posi- 
tive nucleus  surrounded  by  negative  charges  of  electricity,  called 
electrons.     Under  certain  influences  these  electrons  can  pass  pro- 
gressively from  one  atom  to  another,  and  when  their  motion  takes 
place  in  a  single  direction  the  phenomenon  known  as  an  electric 
current  is  observed.     According  to  this  view,  the  effect  produced 
by  such  a  current  is  due  to  a  stream  of  negative  charges  of  elec- 
tricity moving  in  a  single  direction.     Since  different  substances 
differ  markedly  in  electrical  conductivity,  it  is  concluded  from  the 
above  hypothesis  that  they  differ  in  the  ease  with  which  their 
atoms  give  up  electrons.     It  will  be  recalled  -that  according  to  the 
electronic  conception  of  valence  (442),  when  chemical  union  takes 
place  between  two  elements,  an  electron  passes  from  the  positive 
or  metallic  element  to  the  negative  or  non-metallic  element — the 
metallic  element  loses  the  electron.     This  view  is  in  accord  with 
the  fact  that  metals  conduct  the  electric  current,  if  the  hypothesis 
in  regard  to  the  nature  of  the  current  is  accepted.     The  electrons 
present  in  metallic  atoms  can,  with  greater  or  less  ease,  be  set  in 
motion  and  move  progressively  through  the  metal  under  certain 
influences.     Before  these  are  discussed,  however,  it  is  advisable 
to  consider  more  fully  electrical  energy  and  how  it  is  measured. 

563.  Measurement  of  Electrical  Energy. — This  type  of  energy, 
like  all  others,  has  two  factors,  intensity  and  quantity.     A  rough 
analogy  exists  between  the  flow  of  water  through  a  pipe  and  the 
flow  of  an  electric  current  through  a  wire.     The  mechanical  energy 
of  the  moving  water  is  determined  by  the  amount  of  water  in 
motion  and  the  pressure  on  it  which  makes  it  move.     If  a  turbine 
is  driven  by  the  water,  the  amount  of  energy  which  can  be  trans- 
formed into  work  depends  on  these  factors;    it  is  for  this  reason 
that  in  the  development  of  water-power  a  source  is  sought  which 
yields  a  large  amount  of  water — the  quantity  factor — where  the 
fall  is  as  great  as  possible,  in  order  that  a  high  pressure  can  be 
obtained — the  intensity  factor. 

In  an  analogous  way,  when  an  electric  current  flows  through  a 
wire,  the  energy  is  determined  by  the  quantity  that  passes  and  its 
pressure,  which  is  called  in  this  case  electromotive  force  or  poten- 
tial. The  quantity  factor  of  electrical  energy  is  expressed  in  a 


ELECTROCHEMISTRY  473 

unit  called  a  coulomb  1  and  the  unit  of  the  intensity  factor  is 
the  volt;  the  electrical  energy  is  the  product  of  these  two,  and 
is  expressed  in  joules. 

When  electrical  energy  is  being  produced  or  used  and  the  time 
factor  in  the  transformation  is  considered,  additional  units, 
which  express  the  rate  at  which  the  change  takes  place,  are  found 
to  be  convenient.  The  rate  in  the  case  of  the  quantity  factor 
(coulombs)  is  measured  in  amperes;  1  ampere  is  1  coulomb  per 
second.  The  rate  of  transformation  of  electrical  energy  (joules) 
is  expressed  in  watts;  1  watt  is  1  joule  per  second;  1  kilowatt  is 
1000  watts;  1  kilowatt-hour  is  1  kilowatt  for  1  hour  =  1000  X  60 
X  60  joules.  The  use  of  these  terms  is  comparatively  common; 
the  amount  of  electrical  energy  required  to  run  a  lamp  is  usually 
expressed  in  watts  per  candle-power,  and  the  cost  of  electricity  is 
expressed  as  a  certain  number  of  cents  per  kilowatt.  Electric 
lamps  are  also  marked  with  the  voltage  of  the  current  which  must 
be  used  in  connection  with  them,  for,  as  we  shall  see  later,  the 
heat  developed  when  a  current  passes  through  a  lamp  and,  con- 
sequently, the  temperature  of  the  filament  and  the  light  given  off 
by  it  varies  with  the  electromotive  force  of  the  current.  In 
commercial  work  mechanical  energy,  such  as  that  of  a  steam  engine, 
is  expressed  in  horse-power;  1  horse-power  equals  746  watts. 

Instruments  have  been  devised  which  can  be  used  to  measure 
amperes  and  volts;  it  is  only  necessary  to  include  in  an  electric 
circuit  an  ammeter  and  a  voltmeter  to  be  able  to  read  off  on  the 
scale  of  the  instrument  the  values  of  the  quantities  involved. 
Instruments  are  also  used  which  measure  coulombs  and  watts. 

For  the  values  of  these  quantities  expressed  in  absolute  units 
the  reader  is  referred  to  books  on  physics;  but  an  appreciation  of 
their  significance  can  be  obtained  by  noting  the  values  in  the  case 
of  certain  common  electrical  appliances.  The  current  commer- 
cially furnished  for  lighting  and  power  has  an  electromotive  force 
of  110  or  220  volts.  Electric  arc  lights  are  designed  to  take  from 
10  to  250  amperes  according  to  the  brilliancy  of  the  light  required. 
A  tungsten  incandescent  light  consumes  about  1.2  watts  per 
candle-power.  An  ordinary  dry  cell  has  an  electromotive  force 

1  Coulomb  (1736-1806),  Volta  (1745-1827),  Ampere  (1775-1836),  Ohm 
(1787-1854),  Watt  (1736-1819),  and  Joule  (1818-1889),  were  distinguished 
physicists  who  contributed  much  to  our  knowledge  of  energy  and  electricity. 


474  INORGANIC  CHEMISTRY  FOR  COLLEGES 

of  about  1.6  volts,  and  a  single  cell  in  a  storage  battery  about  2 
volts. 

564.  Methods  of  Producing  a  Current  of  Electricity. — Since  an 
electric  current  possesses  energy,  it  is  necessary  to  transform  some 
other  kind  of  energy  in  order  to  produce  it;   mechanical  energy, 
heat,  or  chemical  energy  can  be  used  for  this  purpose. 

From  Mechanical  Energy. — If  two  substances  are  brought  into 
contact  and  then  separated,  they  will  be  found  to  be  charged  with 
electricity,  provided  they  are  insulated  so  that  the  charges  pro- 
duced cannot  be  lost  by  conduction.  The  charge  produced  by  a 
single  contact  is  excessively  small,  but  if  the  objects  are  brought 
together  many  times,  and  one  of  them  is  discharged  after  each 
contact,  the  charge  developed  on  the  other  can  be  made  appre- 
ciable. So-called  frictional  electricity  is  produced  in  this  way. 

The  usual  method  of  converting  mechanical  energy  into  elec- 
trical energy  is  by  means  of  a  dynamo.  When  a  magnet  is  passed 
through  a  circle  made  of  metallic  wire,  a  current  of  electricity  flows 
through  the  wire  as  long  as  the  magnet  is  in  motion;  and  the 
reverse  of  this  is  true  also;  when  an  electric  current  is  passed 
through  a  coil  of  wire  surrounding  a  magnet  the  latter  is  set  in 
motion.  These  fundamental  facts,  discovered  by  Faraday,  were 
utilized  in  the  invention  of  the  dynamo,  which  converts  the  mechan- 
ical energy  of  moving  magnets  into  an  electric  current,  and  of  the 
electric  motor,  in  which  magnets  are  made  to  move  as  the  result 
of  the  passage  of  an  electric  current  around  them. 

565.  From  Heat  Energy. — When  two  metals  are  brought  into 
contact,  they  each  become  charged  with  electricity,  as  we  have 
seen,  but  no  current  flows — the  electricity  is  static.     An  electric 
current  can  be  set  up,  however,  when  two  metals  are  in  contact, 
if  heat  energy  is  supplied.     This  can  be  done  in  a  simple  way.     If 
two  wires  of  different  metals  are  connected  at  their  two  ends  and 
one  of  the  points  of  contact  is  heated,  a  current  of  electricity  flows 
through  the  wires  as  long  as  the  two  junctions  are  at  different 
temperatures.     The  heat  supplied  furnishes  the  energy  necessary 
to  separate  the  electrons  from  the  metals  and  cause  them  to  move 
in  the  wires.     The  amount  of  heat  energy  changed  into  elec- 
trical energy  increases  with  rise  in  temperature.     This  is  the  result, 
according  to  the  accepted  hypothesis,  of  an  increase  in  the  number 


ELECTROCHEMISTRY  475 

of  electrons  set  in  motion  and  the  greater  intensity  with  which 
they  are  given  off. 

The  proportion  of  the  heat  changed  into  electricity  in  this  way 
is  very  small,  and  the  method  is  not  used  as  a  means  of  making 
electrical  energy.  It  is  utilized,  however,  in  an  important  instru- 
ment for  measuring  high  temperatures,  known  as  the  thermoelectric 
pyrometer,  which  is  much  used  in  chemical  work.  The  temperatures 
to  which  the  instrument  is  to  be  subjected  determines  the  metals 
used  in  the  pyrometer;  if  these  are  high,  one  wire  is  platinum  and 
the  other  an  alloy  of  platinum  and  rhodium  or  iridium,  which 
are  rare  metals  of  the  platinum  group.  Instruments  using  these 
metals  are  durable  but  expensive,  and  since  the  metals  are  so  alike 
chemically  the  electromotive  force  of  the  current  set  up  when  they 
are  heated  is  very  small.  If  lower  temperatures  are  to  be  measured, 
metals  farther  apart  in  the  electromotive  series  can  be  used  and  a 
greater  difference  of  potential  is  obtained.  Combinations  com- 
monly used  are  platinum  and  silver,  and  iron  and  chromel,  which 
is  an  alloy  of  nickel  and  chromium. 

In  constructing  a  pyrometer,  the  ends  of  the  two  wires  used 
are  fused  together  and  the  junction  protected  by  a  tube  of  fused 
silica  or  porcelain,  one  end  of  which  is  closed.  This  end  is 
placed  where  the  temperature  is  to  be  measured.  The  other 
ends  of  the  wires  are  kept  at  a  constant  temperature  in  a  vessel 
packed  with  cotton  wool,  or,  when  extreme  accuracy  is  desired, 
in  melting  ice.  To  these  ends  is  attached  a  millivoltmeter— 
an  instrument  designed  to  read  thousandths  of  a  volt.  The 
pyrometer  is  first  calibrated  by  placing  its  terminal  into  sub- 
stances at  known  temperatures,  such  as  boiling  sulphur  and 
certain  metals  at  their  melting-points.  The  reading  on  the 
voltmeter  is  noted  in  each  case,  and  the  results  plotted,  using  tem- 
peratures and  millivolts  as  co-ordinates.  From  the  curve  drawn 
through  these  points  the  temperature  at  any  reading  on  the  volt- 
meter can  be  read  off.  Pyrometers  of  this  kind  are  of  the  greatest 
value  in  the  determination  of  high  temperatures,  and  as  many 
industrial  operations  are  carried  out  at  these  temperatures  the 
instrument  is  much  used.  The  thermoelectric  pyrometer  was 
invented  by  Le  Chatelier,  who  first  discovered  the  law  of  mobile 
equilibrium;  it  is  frequently  designated  by  his  name. 


476  INORGANIC  CHEMISTRY  FOR  COLLEGES 

566.  From  Chemical  Energy. — If  a  rod  of  zinc  and  a  rod  of 
copper  are  placed  in  a  dilute  solution  of  hydrochloric  acid,  the 
metals  not  being  in  contact,  we  observe  that  the  zinc  reacts  with 
the  acid  and  hydrogen  is  given  off  at  the  surface  of  the  metal; 
the  copper  is  not  affected.  If  the  temperature  of  the  solution  is 
carefully  observed  it  will  be  found  that  it  rises  as  the  reaction 
proceeds — chemical  energy  is  being  transformed  into  heat.  If, 
next,  the  two  rods  are  connected  outside  the  solution  by  a  metallic 
wire,  it  will  be  observed  that  the  hydrogen  is  now  evolved  from 
the  copper  rod  and  not  the  zinc;  and  if  a  proper  instrument  for 
detecting  an  electric  current  is  included  in  the  wire  circuit,  it  will 
be  found  that  a  current  is  flowing  in  it.  The  temperature  change 
in  this  case  is  very  small. 

The  chemical  products  of  the  action  in  both  cases  are  the  same; 
zinc  passes  into  solution  as  zinc  chloride,  hydrogen  is  set  free,  and 
the  copper  is  not  altered.  In  the  first  case  all  the  chemical  energy 
lost  when  the  metal  passed  into  solution  was  converted  into  heat; 
in  the  second  case  a  part  of  the  energy  was  transformed  into  elec- 
tricity. 

The  explanation  offered  of  these  striking  facts  is  as  follows: 
When  a  rod  of  zinc  is  placed  in  water  there  is  a  tendency  for  some 
of  the  atoms  of  the  metal  to  pass  into  solution  as  positive  ions, 
which,  it  will  be  recalled,  are  atoms  which  have  lost  one  or  more 
negative  charges  as  electrons,  and  are,  therefore,  said  to  be  positive. 
When  this  takes  place  the  rod  of  zinc  necessarily  assumes  a  nega- 
tive charge  because  it  holds  the  electrons  which  were  formerly 
associated  with  zinc  atoms  which  passed  into  solution.  There  is 
thus  set  up  a  difference  of  potential  between  the  solution  and  the 
metal;  the  change  soon  comes  to  equilibrium  when  the  zinc  ions 
reach  a  certain  concentration.  The  case  is  different  if  an  acid  is 
present.  The  metallic  zinc  is  in  contact  with  hydrogen  ions — 
hydrogen  atoms  each  of  which  has  lost  one  electron.  Since  zinc 
has  a  greater  tendency  than  hydrogen  to  become  an  ion — that  is, 
lose  negative  electrons — it  gives  up  to  the  hydrogen  ions  the  excess 
of  electrons  which  have  accumulated  on  it  as  the  result  of  the 
passage  of  zinc  as  ions  into  water.  When  the  hydrogen  ions  regain 
electrons  they  pass  into  atoms,  form  molecules,  and  hydrogen  gas 
is  given  off. 
We  see,  as  a  result  of  the  consideration  of  the  reaction  between 


ELECTROCHEMISTRY  477 

a  metal  and  an  acid  from  the  point  of  view  of  the  conception 
of  electrons,  that  whether  or  not  a  given  metal  dissolves  in  an 
acid  is  determined  by  the  relative  ease  with  which  it  and  hydro- 
gen lose  these  negative  charges  of  electricity.  The  transfer  of 
the  electrons  from  the  zinc  to  the  hydrogen  atoms  takes  place  with 
a  liberation  of  energy,  which  appears  as  heat.  As  the  electrons 
left  on  the  metallic  zinc  when  the  atoms  of  the  metal  pass  into 
solution  are  immediately  transferred  to  the  hydrogen  ions,  the 
metal  does  not  assume  an  electric  charge. 

The  explanation  can  be  continued  to  include  the  case  where  a 
rod  of  zinc  in  metallic  contact  with  one  of  copper  is  placed  in  hydro- 
chloric acid.  When  pieces  of  copper  and  zinc  are  brought  together 
in  the  absence  of  a  solution,  the  zinc  loses  electrons  to  the  copper 
and  a  difference  in  potential  is  set  up.  We  have  seen  that  a 
current  can  be  produced  if  the  metals  are  brought  into  contact  at 
two  places  to  make  a  circuit  and  one  of  the  junctions  is  heated; 
zinc,  the  more  positive  element,  loses  electrons  to  copper.  When  a 
rod  of  copper  is  joined  to  one  of  zinc  and  the  free  ends  are  placed 
into  a  solution  of  an  acid,  the  electrons  left  on  the  zinc  when  its 
atoms  pass  into  solution,  immediately  flow  through  the  wire  to  the 
copper.  The  excess  of  negative  electrons  is  now  on  this  metal, 
and  it  is  from  the  copper  that  the  hydrogen  ions  get  the  charges 
necessary  to  change  them  into  atoms;  the  gas  is,  accordingly, 
liberated  under  these  conditions  from  the  copper  electrode.  The 
flow  of  negative  electrons  through  the  wire  from  the  zinc  to  the 
copper  produces  an  electric  current.  The  energy  change  asso- 
ciated with  the  conversion  of  metallic  zinc  into  ions  and  of  hydrogen 
ions  into  the  gas  is  largely  the  change  of  chemical  energy  into 
electricity. 

It  should  be  carefully  noted  at  this  point  that  the  terms  used 
in  connection  with  the  flow  of  an  electric  current  are  somewhat 
confusing.  Before  the  advent  of  the  electronic  conception  of  the 
nature  of  the  current,  it  was  the  common  practice  to  assume  that 
the  electric  current  flowed  in  a  circuit  from  positive  to  negative. 
We  are  now  of  the  opinion  that  the  current  is  produced  as  the 
result  of  the  flow  of  negative  electrons  in  the  reverse  direction. 
The  two  views  are  directly  opposite,  but  the  older  method  of 
expressing  the  direction  of  the  flow  of  the  current  is  still  used  with 
the  understanding  that  it  refers  to  what  is  called  the  positive  cur- 


478  INORGANIC  CHEMISTRY  FOR  COLLEGES 

rent.  In  the  electric  cell  just  considered  the  electrons  flow  in  the 
wire  outside  the  liquid  from  zinc  to  copper  and  this  is,  consequently, 
the  direction  of  flow  of  the  negative  current;  but  in  common  prac- 
tice the  current  is  said  to  flow  in  the  opposite  direction.  It  will 
be  easy  to  avoid  confusion  if  it  is  remembered  that  by  definition 
the  current  is  said  to  flow  always  in  the  direction  opposite  to  that 
of  the  moving  electrons,  and  that  when  the  nature  of  the  current 
is  not  specified  this  so-called  positive  current  is  understood. 
Using  this  method  of  expressing  the  direction  of  the  flow  of  the 
current  in  the  cell  described,  the  current  flows  outside  the  solution 
from  the  copper  pole,  which  is  said  to  be  positive  and  is  called  the 
cathode,  to  the  zinc  pole,  which  is  said  to  be  negative  (the  anode) ; 
in  the  cell  it  flows  from  the  zinc  to  the  copper  and  the  circuit  is 
thus  completed.  It  will  be  well  to  remember  that  in  the  solution 
the  positive  current  always  flows  from  the  metal  that  is  passing 
into  solution  as  positive  ions  to  the  metal  upon  which  the  positive 
ions  are  discharged.  The  direction  of  the  flow  of  the  electrons 
and,  consequently,  the  negative  current  is  the  opposite  of 
this. 

567.  Metallic  Couples. — The  explanation  by  means  of  the 
electronic  hypothesis  of  the  way  in  which  an  electric  current  is  set 
up,  can  be  applied  to  the  case  of  metallic  couples  (545).  It  was 
pointed  out  that  zinc  does  not  react  appreciably  with  water,  but 
that  when  it  is  brought  into  contact  with  copper — the  zinc-copper 
couple — hydrogen  is  evolved.  We  have  just  seen  that  when  zinc 
is  put  into  water  some  of  the  metal  passes  into  solution  as  ions. 
The  hydrogen  liberated  is  deposited  on  the  metal  and  the  reaction 
soon  stops.  If  it  is  in  contact  with  copper,  the  gas  is  given  off 
from  this  metal  and  the  zinc  left  free  to  react  further;  the  reac- 
tion is  slow  because  the  zinc  hydroxide  formed  protects  the  metal 
to  some  extent.  Zinc  does  not  react  appreciably  with  a  solution 
of  sodium  hydroxide,  but  hydrogen  is  given  off  from  the  solution 
if  the  zinc-copper  couple  is  used,  for  under  these  circumstances  the 
zinc  hydroxide  is  dissolved  by  the  alkali  as  soon  as  formed. 

Perfectly  pure  zinc  will  not  dissolve  to  any  extent  in  the 
hydrochloric  acid;  the  metal  soon  becomes  covered  with  hydrogen 
on  its  surface  and  reaction  ceases.  If  it  is  impure  it  contains 
substances  which  immediately  set  up  a  difference  of  potential 
with  the  metal  and  the  gas  is  evolved  from  these  impurities,  and 


ELECTROCHEMISTRY  479 

the  zinc  left  free  to  act.  It  was  for  this  reason  that  a  small  amount 
of  a  solution  of  copper  sulphate  was  added  to  zinc  and  the  acid  in 
one  of  our  earliest  experiments  (47).  The  metal  reacted  with  the 
salt,  copper  was  deposited  in  contact  with  the  zinc,  and  a  couple 
was  formed.  If  we  desire  to  prevent  impure  zinc  from  acting 
with  a  dilute  acid,  the  metal  is  coated  with  mercury.  As  a  result 
an  alloy  of  uniform  composition  is  formed  on  the  surface  of  the 
metal  and  there  is  no  opportunity  to  set  up  local  currents  as  is 
the  case  with  the  impure  metal  alone.  It  is  for  this  reason  that 
zinc  rods  and  plates  used  in  the  electric  cells  are  "  amalgamated  " 
with  mercury;  the  metal  is  not  attacked  by  the  materials  in  the 
solution  when  the  cell  is  not  working. 

The  couple  formed  when  the  surface  of  tinned  iron  is  abraded 
so  that  the  iron  is  exposed  to  the  air  acts  in  a  similar  way;  an 
electric  current  is  set  up  in  the  presence  of  water,  the  more  positive 
element — iron — is  converted  into  oxide,  and  hydrogen  is  given 
off  on  the  tin.  In  the  case  of  galvanized  iron  it  is  the  zinc  which  is 
corroded,  as  it  is  the  more  positive  metal. 

568.  An  Electric  Cell. — The  combination  of  zinc,  copper,  and 
hydrochloric  acid  previously  described  could  not  be  used  conveni- 
ently as  a  source  of  electricity,  because  after  a  time  the  copper 
electrode  would  be  covered  with  bubbles  of  hydrogen,  the  zinc 
would  dissolve  in  the  acid  in  the  ordinary  way,  evolving  hydrogen 
on  its  surface,  and  the  chemical  energy  would  be  transformed  into 
heat.  It  is  evident  that  hydrogen  must  not  be  given  off  at  the 
cathode,  and  that  there  should  not  be  enough  hydrogen  ions  present 
to  interact  with  the  zinc  when  the  cell  is  not  being  used  to  produce 
electricity — that  is,  when  the  circuit  is  open.  These  conditions 
are  brought  about  in  a  number  of  ways.  In  the  Daniell  cell  the 
anode  is  zinc  surrounded  by  a  solution  of  zinc  sulphate  placed  in 
a  cup  made  of  baked  clay  which  has  not  been  glazed  and  is,  there- 
fore, porous;  this  rests  in  a  solution  of  copper  sulphate  in  which 
the  cathode  of  copper  is  placed.  It  is  necessary  to  separate  the 
two  solutions,  because  if  the  copper  sulphate  came  into  contact 
with  the  zinc  a  reaction  would  take  place;  zinc  would  pass  into 
solution  and  copper  would  be  deposited  in  the  former.  The  porous 
cup  prevents  the  mixing  of  the  solutions,  but  as  these  penetrate 
its  walls  it  is  possible  for  the  current  to  flow.  When  the  two  poles 
are  connected  outside  the  cell,  the  zinc  passes  into  solution  in  the 


480 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


way  described  above  at  length  and  metallic  copper  is  deposited 
on  the  copper  electrode. 

The  so-called  gravity  cell  (Fig.  34)  is  an  ingenious  improvement 
on  the  one  just  described,  as  it  avoids  the  use  of  a  porous  cup, 
which  adds  to  the  resistance  of  the  cell  and,  therefore,  reduces  the 
current  obtainable  from  it.  The  -copper  electrode  and  copper 
sulphate  crystals  are  placed  at  the  bottom  of  a  glass  jar  which  is 
nearly  filled  with  water,  and  the  zinc  electrode  is  supported  near 
the  top  of  the  jar  under  the  latter.  When  the  terminals  are  con- 
nected a  current  is  set  up  and  after  some  time  the  cell  is  running 


ZnSO4  Solution 


Cu  S04 

Solution 


Copper 
Sulphate 
Crusta/..- 


Copper 


FIG.  34. 


normally.  What  occurs  then  is  as 
follows:  The  presence  of  the  crys- 
tals of  copper  sulphate  keeps  the 
solution  around  the  copper  electrode 
saturated  with  the  salt,  which  fur- 
nishes the  copper  ions  deposited  on 
this  electrode.  The  sulphate  ions 
move  toward  the  zinc  electrode  and 
form  zinc  sulphate  with  the  metal, 
which  passes  into  solution;  there 
is,  thus,  an  accumulation  of  this 
salt  around  the  anode.  The  two 

salts  in  solution  are  kept  from  mixing  for  a  long  time  by  the  fact 
that  a  saturated  solution  of  copper  sulphate  has  a  greater  density 
than  one  of  zinc  sulphate,  and  stays  at  the  bottom  of  the  cell  if  it 
is  not  disturbed;  it  is  for  this  reason  that  it  is  called  a  gravity  cell. 
In  the  so-called  dry  cell,  which  is  so  extensively  used,  zinc  is  the 
anode,  ammonium  chloride  and  zinc  chloride  in  water  the  elec- 
trolyte, and  powdered  manganese  dioxide  and  carbon  in  contact 
with  a  rod  of  carbon  the  cathode.  When  a  current  is  set  up  in 
the  cell  the  ammonium  ions  liberated  at  the  cathode  break  down 
into  ammonia  and  hydrogen.  The  latter  does  not  accumulate 
on  the  cathode,  but  is  oxidized  to  water  by  the  manganese  dioxide. 
The  ammonia  unites  with  the  zinc  salt  in  solution  and  the  satura- 
tion of  the  latter  with  the  free  gas  is  thus  avoided.  The  cell  is 
constructed  in  such  a  way  that  it  is  very  compact  and  efficient. 
The  container  is  the  zinc  which  serves  as  one  electrode.  This 
is  covered  on  the  inside  with  blotting-paper  in  order  to  separate 
the  metal  from  the  powdered  manganese  dioxide,  which  is  packed 


ELECTROCHEMISTRY  481 

in  tight,  after  a  carbon  rod  to  serve  as  a  contact  has  been  placed 
in  the  center  of  the  cell.  A  solution  of  ammonium  chloride  con- 
taining zinc  chloride  is  next  poured  in  to  fill  the  interstices  between 
the  grains  of  the  manganese  dioxide,  and  the  cell  is  finally  closed 
with  melted  pitch.  The  cathode  of  the  dry  cell  is  manganese 
dioxide;  it  is  in  the  form  of  a  closely  packed  powder  which  nearly 
fills  the  entire  cell  and,  accordingly,  has  a  very  large  surface;  it 
is  separated  from  the  zinc  by  the  thickness  of  a  sheet  of  blotting- 
paper  and,  as  a  result,  the  internal  resistance  of  the  cell  is  small. 
These  conditions  make  it  possible  to  get  a  large  amount  of  current 
from  such  a  cell  in  a  given  time — a  dry  cell  connected  directly 
with  an  ammeter  will  develop  as  high  as  30  amperes.  The  large 
surface  of  the  cathode  prevents  the  accumulation  of  hydrogen  on  it, 
and  the  cell  can  be  used  continuously.  If  a  gas  collects  on  an 
electrode  it  prevents  to  a  greater  or  less  extent  contact  between 
the  solution  and  the  electrode;  under  these  conditions  the  elec- 
trode is  said  to  be  polarized. 

In  the  Leclanche  cell  the  electrolyte  is  ammonium  chloride, 
and  the  cathode  is  a  mixture  of  manganese  dioxide  and  carbon 
that  has  been  pressed  into  the  form  of  a  hollow  cylinder;  the 
anode  is  a  rod  of  zinc.  The  chemical  reactions  involved  are 
the  same  as  in  the  dry  cell,  but  as  the  cathode  presents  a  rela- 
tively small  surface  hydrogen  soon  accumulates  on  it  and  as  polari- 
zation takes  place  the  current  soon  drops  off.  The  cell  can  be  used 
only  intermittently. 

In  the  cells  which  have  been  described,  the  electricity  is 
generated  as  the  result  of  utilizing  the  energy  set  free  when  a  metal 
passes  into  solution  as  positive  ions  and  the  positive  ions  of  another 
substance  lose  their  charges  and  come  out  of  solution  in  the  free 
condition — for  example,  Zn  +  Cu+  — >Zn+"  +  Cu.  For  this 
reason  cells  of  this  type  are  called  replacement  cells.  Combina- 
tion cells  based  upon  the  direct  combination  of  two  elements,  oxi- 
dation cells,  and  concentration  cells  in  which  the  electrodes  are 
the  same  metal  in  contact  with  different  concentrations  of  one  of 
its  salts,  have  all  been  fully  studied  and  are  of  interest;  but  their 
consideration  in  a  very  elementary  account  of  electrochemistry 
is  inadvisable. 

569.  The  Quantity  Factor  of  the  Energy  in  Electrochemical 
Change. — In  the  consideration  of  the  electric  cell  up  to  this  point 


482  INORGANIC  CHEMISTRY  FOR  COLLEGES 

attention  has  been  paid  to  the  way  in  which  the  current  is  pro- 
duced, but  nothing  has  been  said  as  to  the  quantity  of  electricity 
formed  or  its  intensity  factor.  The  relation  between  the  amount 
of  electricity  produced  and  the  weight  of  the  metal  involved  is 
best  determined  by  studying  the  conversion  of  metallic  ions  into 
free  metals  when  a  current  of  electricity  is  passed  through  solu- 
tions of  their  salts.  This  process  is  just  the  reverse  of  that  by 
which  a  current  is  formed — the  metallic  ion  loses  its  positive 
charge  as  the  result  of  the  taking  up  of  electrons  and  is  thereby 
changed  to  the  metallic  condition.  These  changes  can  be  indi- 
cated as  follows  where  ©  represents  an  electron — a  negative 
charge:  Zn+  +  2  ©  ;=±  Zn.  Read  from  left  to  right  the  equa- 
tion indicates  that  a  zinc  ion  takes  up  two  negative  charges  and  is 
thus  changed  to  the  free  metal.  This  is  what  happens  when 
zinc  is  deposited  by  electricity  from  a  solution  of  one  of  its  salts. 
Read  in  the  opposite  direction  the  equation  means  that  the  metal 
loses  two  electrons  and  becomes  an  ion;  the  change  indicated  is 
the  production  of  electricity. 

The  action  of  the  electric  current  on  solutions  was  studied  by 
Faraday  in  1834,  and  important  facts  were  discovered.  It  was 
shown,  first,  that  the  quantity  of  any  one  substance  decomposed 
by  the  electric  current  is  proportional  to  the  quantity  of  elec- 
tricity passed  through  its  solution,  and,  second,  that  the  quanti- 
ties of  two  or  more  substances  liberated  by  equal  quantities  of 
electricity  are  proportional  to  the  chemical  equivalents  of  these 
substances.  This  statement  is  known  as  Faraday's  law.  By 
chemical  equivalent  is  meant  the  weight  equivalent  to  1.008  grams 
of  hydrogen.  This  is  1  gram-atomic-weight  of  a  univalent  ele- 
ment, J  gram-atomic-weight  of  a  bivalent  element,  etc.  This 
law  is  of  fundamental  importance  and  is  the  basis  for  all  calcu- 
lations involving  the  relationship  between  quantities  of  elec- 
tricity, measured  in  coulombs,  and  weights  of  substances  in  grams. 
It  was  found  by  experiment  that  96,500  coulombs  are  required  to 
liberate  1  gram-atomic-weight.  Accordingly,  this  quantity  of 
electricity  will  liberate  1.008  grams  of  hydrogen,  35.46  of  chlorine, 
23.00  grams  of  sodium,  65.37  -^  2  =  32.68  grams  of  zinc,  27.1  -T- 
3  =  9.03  grams  of  aluminium,  etc.  Since  reference  is  often  made 
to  the  quantity  of  electricity  which  liberates  one  equivalent  of  an 


ELECTROCHEMISTRY  483 

element,  96,500  coulombs,  it  has  been  given  a  special  name  and  is 
called  a  faraday. 

Under  carefully  controlled  conditions  the  deposition  of  metals 
by  the  electric  current  from  solutions  of  their  salts  can  be  carried 
out  with  great  accuracy,  and  this  method  is  used  in  an  instrument 
called  a  coulometer  to  measure  quantities  of  electricity.  It  is  simple 
in  construction.  Two  pieces  of  copper  placed  in  a  solution  of  a 
copper  salt  are  connected  with  the  poles  of  the  source  of  electricity 
to  be  measured.  As  the  current  passes  through  the  instrument  the 
metal  is  dissolved  at  one  pole  and  deposited  at  the  other.  The 
increase  in  weight  of  the  pole  connected  with  the  anode  is  a  measure 
of  the  current  passing  through  the  instrument;  for  each  gram- 
atomic-weight  of  copper  deposited,  63.57  grams,  2  X  96,500 
coulombs  of  electricity  passed.  If  the  time  required  for  the  depo- 
sition is  noted  the  amperes  can  be  calculated. 

We  have  just  seen  the  relationship  between  the  quantities  of 
electricity  and  weights  of  substances  which  are  involved  in  the 
change  represented  by  the  equation  Zn+  +  +  2  ©  =  Zn.  If  the  reac- 
tion is  carried  out  in  the  reverse  direction  and  the  zinc  is  changed 
into  ions  under  the  conditions  which  permit  the  formation  of  an 
electric  current,  as  in  the  gravity  battery,  the  quantity  of  elec- 
tricity produced  when  a  given  weight  of  the  metal  dissolves,  is 
just  equal  to  that  required  to  deposit  the  same  weight  of  metal  from 
its  ions.  Thus,  1  gram-atomic-weight  of  a  univalent  element  in 
passing  into  the  ionic  condition  will  liberate  96,500  coulombs; 
likewise  one-half  of  a  gram-atomic-weight  of  a  bivalent  element 
will  liberate  this  quantity,  etc.  In  brief,  each  valence  of  an 
element  when  it  appears  or  disappears  involves  a  transfer  of  96,500 
coulombs,  when  we  consider  gram-atomic-weights  in  all  cases. 

We  are  now  in  a  position  to  understand  the  quantity  of  elec- 
tricity that  can  be  obtained  from  a  given  cell.  This  is  deter- 
mined by  the  weight  of  the  element  which  is  dissolved;  each  gram- 
equivalent  furnishes  96,500  coulombs.  The  time-factor  involved 
is  an  important  one;  two  cells  may  furnish  the  same  amount  of 
electricity,  but  one  may  do  this  in  a  few  minutes  and  the  other 
require  hours.  The  rate  at  which  the  current  is  furnished  is  deter- 
mined, of  course,  by  the  rate  at  which  the  metal  dissolves,  and  this 
is  proportional  to  the  surface  of  the  latter  exposed  in  the  cell.  If 


484  INORGANIC  CHEMISTRY  FOR  COLLEGES 

the  zinc  poles  of  two  cells  of  the  same  size  and  construction  are 
connected  and  the  copper  poles  are  likewise  connected,  we  can 
obtain  from  the  combination  twice  the  quantity  of  electricity  in  a 
given  time  that  can  be  obtained  from  one  of  the  cells;  the  com- 
bination delivers  twice  as  many  amperes  as  a  single  cell.  By 
joining  the  two  zinc  poles  we  have  doubled  the  area  of  the  latter 
and  the  amount  of  the  metal  dissolved  in  a  given  time  is  doubled. 
Cells  joined  in  this  way  are  said  to  be  connected  in  parallel. 

570.  The  Intensity  Factor  of  the  Energy  Transformed  in 
Electrochemical  Change. — The  fact  has  been  mentioned  a  number 
of  times  that  when  two  different  substances  are  brought  into  con- 
tact a  difference  of  potential  is  set  up  immediately  as  the  result 
of  the  transfer  of  electrons  from  one  to  the  other.  This  occurs 
when  a  metal  is  placed  in  a  solution  of  one  of  its  salts;  in  the  case 
of  zinc,  for  example,  we  have  Zn  —  2  ©  ^±  Zn++.  Some  zinc  atoms 
lose  electrons  and  pass  into  solution  as  ions  charged  positively  in 
respect  to  the  metal,  because  the  latter  holds  an  excess  of  elec- 
trons and  is,  therefore,  negatively  charged.  The  ease  with  which 
this  change  takes  place  is  determined  in  the  case  of  any  metal  by 
the  number  of  ions  present  in  the  solution;  the  more  ions  present 
the  less  the  tendency  of  the  metal  to  form  ions,  and  the  reverse  is 
true.  As  a  consequence,  in  studying  the  phenomenon  quanti- 
tatively the  concentration  of  the  ions  in  the  solution  must  be 
known. 

The  metals  differ  from  one  another  in  the  strength  of  the 
tendency  they  exhibit  to  form  ions.  Sodium,  for  example,  passes 
readily  into  solution  as  ions  when  brought  into  contact  with  water; 
calcium  forms  ions  less  readily;  in  the  case  of  zinc  the  tendency  is 
much  smaller.  There  is  an  analogy  between  the  tendency  of 
liquids  to  pass  into  vapor  and  the  tendency  of  metals  to  pass 
into  the  forms  of  ions.  In  the  case  of  the  former  the  tendency  is 
measured  by  the  pressure  of  the  vapor  produced  from  the  liquids; 
in  the  case  of  the  latter  the  tendency  is  measured  by  what  is  called 
the  electrolytic  solution  pressure.  This  pressure  is  the  cause  of  the 
difference  of  potential  set  up  when  a  metal  is  brought  into  contact 
with  water  or  a  solution  of  one  of  its  salts.  In  order  to  measure 
the  solution  pressure  of  a  metal  it  is  only  necessary,  therefore,  to 
determine  this  difference  in  potential. 

When  a  voltmeter  is  connected  with  the  two  poles  of  a  cell,  the 


ELECTROCHEMISTRY  483 

instrument  measures  the  difference  between  the  potential  of  one 
pole  and  that  of  the  other;  the  electromotive  force  observed  is  the 
difference  between  two  potentials  In  the  case  of  the  Daniell  cell, 
for  example,  the  potential  of  the  zinc  pole  is  produced  as  the  result 
of  the  passage  of  metallic  zinc  into  ions,  and  that  of  the  copper  pole 
as  the  result  of  the  change  of  copper  ions  to  the  metal.  There  is  a 
difference  in  potential  between  the  zinc  pole  and  the  solution  of 
ions,  and  one  between  the  copper  pole  and  its  ions.  In  order  to 
distinguish  between  a  difference  of  potential  set  up  between  a  metal 
and  a  solution  of  its  salt,  and  one  set  up  between  the  two  poles  of  a 
cell,  the  former  type  is  called  a  single  potential.  The  name  is  a 
satisfactory  one  because  a  single  metal  is  involved  in  setting  up  the 
difference  in  potential,  whereas  the  difference  in  potential  of 
the  two  poles  of  a  battery  results  from  two  potentials  set  up  by 
two  metals. 

When  some  metals  are  placed  in  normal  solutions  of  their  salts 
there  is  a  greater  tendency  for  the  metal  to  pass  into  solution  than 
for  the  ions  to  deposit  as  free  metal;  with  certain  metals  the 
reverse  is  true.  In  the  first  case  the  metal  becomes  charged  nega- 
tively and  the  solution  positively,  because  some  atoms  pass  into 
solution  without  electrons  to  form  the  ions  and  the  excess  of  elec- 
trons left  in  the  metal  charge  it  negatively;  the  solution  is  posi- 
tively charged  because  it  contains  an  excess  of  positive  ions.  In 
the  second  case  the  metal  is  charged  positively  and  the  solution 
negatively;  this  results  from  the  fact  that  when  the  positively 
charged  ions  deposited  on  the  metal  they  rendered  the  latter 
positive. 

The  metals  differ  not  only  in  regard  to  the  positive  or  negative 
nature  of  the  charge  they  assume  when  placed  in  contact  with 
their  solutions,  but  in  the  intensity  of  the  charge;  the  difference  in 
potential  produced  between  a  metal  and  a  normal  solution  of  its 
ions  is  definite  and  characteristic  of  the  metal.  The  values  for  the 
single  potentials  produced  in  this  way  have  been  determined  for 
all  the  metals.  When  the  metal  reacts  with  water  at  ordinary 
temperatures  its  single  potential  cannot  be  determined,  but  can 
be  calculated  from  data  obtained  in  other  ways.  The  values  of 
the  single  potentials  set  up  when  the  metals  are  placed  in  normal 
solutions  of  their  salts  are  given  in  the  following  table.  This 
arrangement  of  the  metals  is  known  as  the  electromotive  series. 


Fe(Fe++)., 

,  +  0.14 

Sb  

—  0  8? 

Co. 

—  0  05 

Hg  (Hg+) 

—  1  02 

Ni  

—  0  05 

Ag 

—  1  08 

Sn(Sn++)., 

,  —  0.13 

Pd  

.  —  11? 

Pb  (Pb++). 

—  0.15 

Pt  

—  1  2? 

H2 

—  0  28 

Au  (Au 

')  —  1  4? 

Cu  (Cu  ++). 

-  0.60 

486  INORGANIC  CHEMISTRY  FOR  COLLEGES 

ELECTROMOTIVE  SERIES  OF  THE  METALS 

The  figure  after  each  metal  is  the  potential  in  volts  of  a  normal  solution 
of  a  salt  of  the  metal  in  contact  with  the  metal.  The  figures  in  parentheses 
have  not  been  determined  directly;  the  values  which  are  followed  by  an 
interrogation  point  are  doubtful. 

K (+2.6)        Cd +  0.15          Bi 

Na (+2.4) 

Ba (+2.4) 

Sr (+2.3) 

Ca (+1.9) 

Mg +1.5? 

Al +1.0? 

Mn +0.8 

Zn +0.5          As -0.6? 

The  signs  +  and  —  in  the  table  refer  to  the  charge  on  the  solu- 
tion. Down  to  cobalt  the  solutions  are  positively  charged;  this 
means  that  there  is  a  greater  tendency  for  the  metal  to  form  ions 
than  for  the  ions  to  pass  into  the  metallic  condition.  From  cobalt 
to  gold  the  greater  tendency  is  in  the  reverse  direction,  and  the 
solutions  become  negatively  charged.  In  the  case  of  hydrogen  an 
electrode  is  formed  by  keeping  saturated  with  the  gas  finely 
divided  platinum  supported  on  platinum  foil. 

671.  In  the  above  table  the  single  potential  of  the  hydrogen  electrode 
is  taken  as  —0.28  volt.  The  single  potentials  of  the  other  elements  have 
been  determined  by  comparing  them  with  that  of  hydrogen.  No  method 
has  been  devised  of  determining  in  an  entirely  satisfactory  manner  the  single 
potential  set  up  between  an  element  and  its  ions.  Consequently,  there  is 
doubt  as  to  the  correctness  of  the  value  assigned  to  hydrogen.  On  account 
of  this  fact,  many  chemists  take  the  hydrogen  electrode  as  the  standard,  and 
assign  to  its  single  potential  the  value  0.  The  values  of  the  single  potentials 
on  this  basis  can  be  obtained  by  adding  0.28  to  the  numbers  given  in  the 
table.  The  elements  fall  in  the  same  order  whatever  standard  is  used;  and 
it  is  this  order  which  is  of  importance  in  interpreting  the  properties  of  the 
elements  and  their  compounds. 

572.  The  electromotive  series  of  the  metals  has  been  used 
repeatedly  throughout  this  book  in  interpreting  the  behavior  of 
metals  in  a  comparative  way.  We  see  now  how  it  has  been 
arrived  at.  The  metals  are  arranged  in  the  order  of  their  ten- 
dency to  pass  into  solutions  as  ions  as  measured  by  the  intensity 
factor  of  the  energy  developed  when  such  a  change  takes  place. 
We  have  seen  that  the  quantity  factor  of  the  electrical  energy  set 


ELECTROCHEMISTRY  487 

free  in  this  way  is  determined  by  the  valence  of  the  ion.  When 
1  gram-atomic-weight  of  magnesium,  which  has  the  valence  2, 
passes  into  solution  as  a  magnesium  ion,  2X96,500  coulombs  of 
electricity  are  produced;  the  same  quantity  is  set  free  when  1 
gram-atomic-weight  of  iron  passes  into  bivalent  ions.  In  the 
first  case  1.5  volts  are  developed,  and  in  the  second  case  but  0.14 
volt.  Since  electrical  energy  is  the  product  of  its  quantity  and 
intensity  factors,  the  chemical  energy  transformed  into  electrical 
energy  in  the  case  of  magnesium  is  2  X  96,500  X  1.5  =  289,500 
joules  and  in  the  case  of  iron  2  X  96,500  X  0.14  =  27,020  joules. 
These  facts  make  clear  the  difference  in  the  energy  content  and 
activity  of  the  two  metals. 

The  values  of  the  potentials  given  in  the  table  are  those  obtained 
when  normal  solutions  of  salts  are  in  contact  with  the  metals. 
At  other  concentrations  the  values  are  different.  For  example, 
in  the  reaction  Zn  -  2©  ^±Zn++,  which  is  reversible,  an  equi- 
librium is  set  up  as  the  result  of  the  opposing  tendency  of  the  metal 
to  pass  into  ions,  and  the  tendency  of  the  ions  to  deposit  as  metal. 
Increase  in  the  number  of  metallic  ions  would,  accordingly,  bring 
about  a  new  state  of  equilibrium,  and  the  difference  in  potential 
set  up  between  the  metal  and  the  solution  would  be  less  than 
before.  In  comparing  the  activity  of  two  metals  as  measured 
by  the  potentials  set  up,  false  conclusions  maybe  arrived  at  if 
the  concentrations  of  the  ions  of  the  metals  are  widely  different. 
When  a  metal  above  hydrogen  in  the  electromotive  series  reacts 
with  an  acid  to  form  hydrogen,  it  dissolves  because  its  tendency  to 
form  ions  is  greater  than  that  of  hydrogen ;  the  metal  passes  into 
solution  and  the  hydrogen  ions  are  discharged.  As  the  metal 
dissolves  and  its  ions  increase  in  concentration,  the  potential 
becomes  less  and  less  but,  in  general,  its  value  remains  greater 
than  that  of  hydrogen,  and  the  gas  continues  to  be  evolved. 

573.  In  the  case  of  a  metal  below  hydrogen,  the  gas  is  not 
ordinarily  set  free  when  it  is  placed  in  a  solution  of  an  acid, 
because  hydrogen  has  a  greater  tendency  to  form  ions  than  the 
metal  under  the  conditions  which  exist.  In  this  case  there  is  an 
equilibrium  set  up,  however,  between  the  elements  and  their  ions, 
as  before,  and  if  the  metallic  ions  are  removed  as  fast  as  they  are 
formed  the  tendency  of  the  metal  to  form  ions  must  increase ;  if  it 
becomes  greater  than  that  of  hydrogen  the  metal  passes  into  solu- 


488  INORGANIC  CHEMISTRY  FOR  COLLEGES 

tion.  This  is  what  occurs  when  mercury  is  treated  with  concen- 
trated hydrobromic  acid,  and  copper  with  concentrated  hydro- 
chloric acid.  In  both  cases  compounds  are  formed  between  the 
acids  and  the  halides  produced,  which  yield  an  exceedingly  small 
concentration  of  metallic  ions. 

574.  Electromotive  Series  of  Non-metallic  Ions. — The  single 
potentials  set  up  between  certain  non-metallic  elements  and  their 
ions  have  been  determined;  some  of  these  are  given  in  the  follow- 
ing table: 

I2(I~)    -0.81  Cl2(Cr)    -1.64 

Br2(Br~)    -1.36  O2(OH-)    -1.94 

02(0~-)-1.49 

When  chlorine  is  in  contact  with  water  some  of  the  element  passes 
into  solution  as  Cl~  ions;  as  a  result,  the  solution  becomes  nega- 
tively charged  in  respect  to  the  free  element.  The  sign  before  the 
number  expressing  the  difference  in  potential  is,  therefore,  nega- 
tive, since  in  this  table,  as  well  as  in  that  of  the  electromotive  series 
of  metals,  the  sign  refers  to  the  charge  on  the  solution.  It  indicates 
in  the  case  of  all  the  elements  listed  above  a  tendency  to  pass  into 
solution  as  negative  ions.  The  numerical  value  of  the  difference 
in  potential,  disregarding  the  sign,  is  a  measure  of  the  tendency 
for  ionization  to  take  place — the  larger  the  number  the  greater 
this  tendency.  It  is  evident  from  an  examination  of  the  figures 
why  free  chlorine  displaces  iodine  atoms  from  solutions: 

2K+  +  21-  +  C12  =  2K+  +  I2  +  2Cr 

675.  The  Electromotive  Force  of  Electric  Cells.— When  metals 
and  solutions  of  salts  are  combined  to  make  a  cell,  the  electro- 
motive force  of  the  combination  is  the  result  of  two  separate  elec- 
tromotive forces  produced  in  the  cell;  one  is  set  up  between  the 
anode  and  the  ions  of  the  metal  of  which  it  is  composed,  and  the 
other  between  the  cathode  and  its  ions.  From  the  values  given  in 
the  table  on  page  486  it  is  possible  to  calculate  the  electromotive 
force  of  a  cell  composed  of  two  metals,  each  in  contact  with  a 
normal  solution  of  one  of  its  salts.  It  will  be  recalled  that  the 
signs  given  in  the  table  refer  to  the  charge  on  the  solutions  and 
that  those  on  the  metals  have  the  opposite  sign.  In  the  Daniell 
cell,  for  example,  copper  is  charged  positively  and  its  potential 


ELECTROCHEMISTRY  489 

with  reference  to  copper  sulphate  is  +0.60  volt;  zinc  is  charged 
negatively  and  its  potential  is  —0.5  volt.  The  difference  in  poten- 
tial between  the  two  metals  is  evidently  the  difference  between 
these  two  values,  0.60  —  (—0.5)  =  1.1  volts.  If  a  cell  is  made 
up  of  iron  and  zinc  the  difference  in  potential  is  —0.14  —  (  —  0.5)  = 
0.36  volt. 

In  order  to  get  as  high  an  electromotive  force  as  possible 
elements  widely  separated  in  the  electromotive  series  are  selected. 
It  is  found  in  practice  that  zinc  on  account  of  its  cheapness  is  the 
best  metal  to  use  high  in  the  series.  Either  copper  or  carbon 
with  manganese  dioxide  are  the  most  available  substances  from 
which  to  construct  cathodes.  When  manganese  dioxide  is  used, 
the  reaction  at  the  cathode  which  produces  an  electromotive  force 
is  the  reduction  of  MnO2  to  M^Oa,  the  oxygen  reacting  with 
hydrogen  to  form  OH~  ions.  The  E.M.F.  of  a  dry  cell  is  approx- 
imately 1.6  volts. 

In  but  a  few  cases  it  is  possible  to  construct  a  cell  in  which 
all  the  chemical  energy  set  free  is  converted  into  electrical  energy; 
in  most  cases  there  is  a  heat  change,  and  when  this  occurs  the  elec- 
tromotive force  is  not  that  calculated  in  the  way  indicated 
above.  This  difference  in  the  case  of  cells  commonly  used  is 
small,  however,  and  the  change  in  potential  due  to  differences  in 
concentration  of  the  ions  is  likewise  relatively  small  for  com- 
paratively great  changes  in  concentration.  The  calculated  values 
approximate,  therefore,  the  observed  ones. 

576.  Electrolytic  Conduction. — It  has  been  pointed  out  that 
when  an  electric  current  flows  through  a  wire  there  is  no  chemical 
change  in  the  material  of  which  the  wire  is  made;  as  the  current 
passes  an  atom  gives  up  an  electron  to  the  next  atom  and  receives 
another  in  turn,  which  it  passes  on;  a  constant  stream  is  thus 
produced  if  a  supply  of  electrons  is  furnished  at  one  end  of  the 
wire  and  they  are  removed  continuously  at  the  other  end.  The 
conduction  in  a  solution  where  the  current  is  carried  by  the  ions 
is  pictured  differently.  In  this  case  where  the  conduction  is  said 
to  be  electrolytic,  the  electrolyte,  which  is  the  acid,  base,  or  salt  in 
solution,  undergoes  decomposition  at  the  electrodes.  The  nega- 
tive ions  give  up  electrons  at  one  pole  and  the  positive  ions  take 
up  electrons  at  the  other.  In  the  case  of  hydrochloric  acid,  for 
example,  each  chlorine  ion  loses  an  electron  and  is  thereby  con- 


490 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


verted   into   atomic   chlorine,   which   changes   immediately   into 
chlorine  gas;    likewise  the  hydrogen  ions  take  up  an  electrons 

x  and  become  hydrogen  atoms  and 

then  hydrogen  gas.  When  the 
current  has  passed  a  short  time 
the  ions  in  the  immediate  vicinity 
of  the  electrode  have  been  dis- 
the 


,  Sit 


FIG.  35. 


Fre&  charged,  but  tne  process  con- 
tinues because  the  charged  elec- 
trodes attract  to  themselves  the 
charged  ions  in  the  solution. 
There  is  thus  set  up  a  continuous 
stream  of  negative  ions  toward 
the  pole  charged  positively  and 
one  of  positive  ions  toward  the 
negative  pole;  the  positive  ions 
travel  in  the  direction  of  the  flow 

of  the  positive  current  and  the  negative  ions  in  the  direction  of 
the  flow  of  the  electrons  or  negative  current.  These  conditions 
are  shown  in  the  accompanying  diagram  (Fig.  35). 

577.  Storage  Batteries. — In  the  type  of  cell  which  has  been 
described,  electrical  energy  is  produced  from  chemical  energy  as 
the  result  of  the  change  of  a  metal  into  ions  and  the  reverse.  The 
current  in  the  Daniell  cell  is  set  up  as  the  result  of  changes  which 
may  be  indicated  thus : 

Zn  =  Zn++  +  2  ©  and  Cu++  +  2  ©  =  Cu 

Zinc  passes  into  solution  at  one  pole  and  copper  is  deposited  at 
the  other,  both  changes  liberating  energy  which  sets  up  a  flow  of 
electrons  from  the  zinc  pole,  where  they  were  set  free,  through  the 
connecting  wire  to  the  copper  pole,  where  they  were  taken  up,  and, 
as  a  consequence,  an  electric  current  is  produced.  Both  reac- 
tions are  reversible,  and  if  electrons  are  furnished  from  outside 
the  cell  and  made  to  flow  in  a  direction  opposite  to  that  in  which 
they  travel  when  the  cell  is  producing  electricity,  the  chemical 
changes  effected  will  be  reversed.  When  the  cell  is  producing 
electricity  the  reaction  at  the  zinc  pole  is  Zn  =  Zn++  +  2  ©; 
when  a  supply  of  electrons  is  furnished  by  an  outside  current  the 
reaction  is  Zn++  +  2  ©  =  Zn,  and  a  similar  reversal  takes  place/  at 


ELECTROCHEMISTRY  491 

the  copper  pole.  It  is  evident,  therefore,  that  we  can  obtain 
electricity  from  the  Daniel!  cell,  or,  by  passing  electricity  into  it, 
bring  about  chemical  changes  which  lead  to  the  separation  of  sub- 
stances from  which  electricity  can  be  produced;  in  the  latter 
case  electrical  energy  is  transformed  into  chemical  energy,  which 
later  can  be  reconverted  into  electricity.  When  the  Daniell  cell 
is  used  in  the  first  way  it  is  said  to  be  a  primary  cell;  when  used 
in  the  second  way  it  is  said  to  be  a  secondary  or  storage  cell. 

For  many  reasons  the  Daniell  cell  is  not  economical  when  used 
as  a  storage  cell,  and  other  combinations  are  employed.  In  all 
cases,  substances  are  selected  which  react  reversibly  without  appre- 
ciable loss  of  energy  as  heat,  and  without  the  production  of  gases 
which  escape  from  the  cell.  There  are  at  present  two  important 
kinds  of  storage  cells. 

578.  In  the  lead  accumulator,  when  it  is  charged,  one  pole  is 
lead  dioxide,  PbO2,  the  other  metallic  lead,  and  the  electrolyte 
is  sulphuric  acid.  When  the  poles  are  connected  so  that  a  current 
can  flow,  lead  sulphate  is  formed  at  both  the  anode  and  cathode. 
Metallic  lead  reacts  with  the  sulphate  ion  to  form  lead  sulphate, 

Pb  +  SO4~  ~  =  PbSO4  +  20 

and  the  electrons  on  the  negative  ion  are  set  free.  The  hydrogen 
ions  at  the  other  pole  take  up  electrons  and  are  thus  converted 
into  atoms  which  reduce  the  dioxide  to  oxide;  and  the  latter  is 
converted  by  the  acid  present  into  lead  sulphate  and  water : 

2H+  +  20=  [2H] 
[2H]  +  PbO2  =  H2O  +  [PbO] 
[PbO]  +  H2S04  =  PbSO4  +  H20 
2H+  +  PbO2  +  H2SO4  +  20  =  2H2O  +  PbSO4 

The  current  is,  thus,  produced  as  the  result  of  the  liberation  of  2 
electrons  from  the  SO4  radical  at  one  pole  and  the  transfer  of 
2  electrons  to  2H+  ions  at  the  other.  The  complete  equation  for 
the  reaction  which  takes  place  during  the  discharge  of  the  cell  is, 
accordingly, 

PbO2  +  Pb  +  2H2SO4  =  2PbS04  +  2H20  +  Energy 


492  INORGANIC  CHEMISTRY  FOR  COLLEGES 

When  a  current  is  passed  into  the  cell  to  charge  it,  the  reaction 
at  one  pole  is 

PbSO4 


At  the  other  pole  an  SO4  ion  is  taken  up  to  form  a  persulphate 
which  is  immediately  hydrolyzed  by  water: 

PbSO4  +  SO4~~  =  [Pb(SO4)2J  +  20 
[Pb(SO4)2]  +  2H2O  =  PbO2  +  2H2SO4 
PbSO4  +  SO4—  +  2H2O  =  PbO2  +  2H2SO4  +  20 

The  summation  of  the  reactions  at  the  two  poles  leads  to  an 
equation  identical  with  that  which  represents  the  discharge  of  the 
cell,  except  that  it  is  written  in  the  reverse  direction  : 

charge 

Energy  +  2PbS04  +  2H2O  ^±  PbO2  +  Pb  +  2H2SO4 

discharge 

It  is  not  difficult  to  remember  in  which  direction  the  equation  is  to 
be  read  to  express  either  discharge  or  charge,  when  it  is  recalled 
that  the  formation  of  a  salt  and  water  from  an  acid  and  a  metal 
or  an  oxide  is  associated  with  the  setting  free  of  energy.  Under 
ordinary  circumstances  this  energy  appears  as  heat,  but  by  the 
arrangement  used  in  a  storage  cell  it  is  given  off  in  the  form  of  an 
electric  current.  The  cause  of  the  setting  up  of  the  current  can 
be  expressed  very  simply;  it  is  produced  as  the  result  of  the  change 
of  one  atom  of  metallic  lead,  which  has  0  valence,  Pb°,  and  of  1 
atom  of  lead  as  the  dioxide  with  the  valence  4,  PbIV,  into  2  atoms  of 
lead  in  the  form  of  the  sulphate  in  which  the  metal  has  the 
valence  2,Pb°  +  PbIV  =  2Pbn.  The  metallic  lead  loses  2  electrons, 
Pb°  =  Pb+  +20,  which  convert  the  quadrivalent  lead  to 
bivalent  lead,  2  0  +  Pb  +  +  +  +  =  Pb  +  +  ;  it  is  the  transfer  of  these 
electrons  in  the  circuit  which  causes  the  current. 

The  poles  of  the  lead  accumulator  are  made  in  the  form  of  flat 
grids  of  lead  the  interstices  of  which  contain  the  active  materials  of 
the  cell,  finely  divided  lead  in  one  case  and  lead  dioxide  in  the  other. 
The  electromotive  force  produced  by  each  cell  is  about  2  volts. 
When  a  cell  is  in  the  charged  condition  the  solution  contains  sul- 
phuric acid;  as  it  is  discharged  lead  sulphate  is  deposited  on  both 
poles  and  water  is  formed.  It  is  possible,  therefore,  by  an  observa- 


ELECTROCHEMISTRY  493 

tion  of  the  specific  gravity  of  the  solution,  which  is  a  measure  of 
the  amount  of  acid  present,  to  tell  to  what  extent  the  cell  has  been 
discharged. 

579.  The  lead  accumulator  is  very  heavy,  since  it  is  made  of 
lead  and  involves  the  use  of  sulphuric  acid,  and,  consequently, 
many  attempts  have  been  made  to  utilize  other  reactions  in  the 
construction  of  storage  cells.     The  most  successful  of  these  is 
known  as  the  Edison  cell,  in  which  the  poles,  when  charged,  are 
nickelic  oxide,  Ni2O3,  and  iron.     The  electrolyte  is  a  solution  of 
potassium    hydroxide.     When    the    cell    discharges    the    nickelic 
oxide  is  changed  to  nickelous  hydroxide,  Ni(OH)2>  and  the  metallic 
iron  to  ferrous  hydroxide : 

discharge 

Fe  +  Ni2O3  +  3H2O  ^±  Fe(OH)2  +  2Ni(OH)2  +  Energy 

charge 

The  changes  in  valence  which  cause  the  current  are  represented  as 
follows : 

discharge 

Fe°  +  2Nim  ^±  Fe11  +  2Nin  +  Energy 

charge 

Iron  is  above  nickel  in  the  electromotive  series  and,  as  a  conse- 
quence, the  energy  set  free  when  it  is  oxidized,  Fe°  — *  Fe11,  is 
greater  than  that  required  to  reduce  an  equivalent  amount  of 
nickel,  2Nini  — *  2Nin;  the  conversion  in  this  direction  thus 
liberates  energy  which  appears  in  the  form  of  electricity. 

580.  Uses   of   the   Electric   Current   in   Chemistry,     (a)  For 
Heating. — All  substances  offer  more  or  less  resistance  to  the  flow 
of  an  electric  current  through  them,  and,  as  a  consequence,  a  part 
of  the  electrical  energy  is  changed  into  heat.     When  it  is  desired 
to  avoid  loss  of  energy  in  this  way,  as  in  the  conduction  of  elec- 
trical energy  from  one  place  to  another,  a  conductor  is  chosen 
which  has  a  low  resistance;    copper  is  ordinarily  used  for  this 
purpose  (see  table,  page  443).     When,  however,  it  is  desired  to 
convert  electrical  energy  into  heat,  a  conductor  which  has  a  high 
resistance  is  selected,  and  for  this  reason  alloys,  carbon,  and  non- 
metallic  substances,  are  commonly  used.     The  heat  generated  in  a 
circuit  is  equal  to  the  square  of  the  current  multiplied  by  the 
resistance.     When  the  current  is  1  ampere  and  the  resistance  is 
1  ohm  the  heat  developed  is  0.2388  calorie  (  =  1  joule  per  second). 


494  INORGANIC  CHEMISTRY  FOR  COLLEGES 

When  comparatively  small  quantities  of  heat  are  to  be  devel- 
oped, a  current  is  sent  through  a  wire  in  order  to  deliver  the  heat 
where  it  is  required.  For  small  furnaces  tubes  of  fused  silica  or 
porcelain  are  wound  with  a  wire  made  of  an  alloy  which  has  a 
high  resistance  and  a  high  melting-point,  and  does  not  oxidize 
rapidly  in  air  at  high  temperatures.  An  alloy  containing  65  per 
cent  chromium  and  35  per  cent  nickel  is  well  adapted  to  this  pur- 
pose^commercial  "  nichrome "  or  "  chromel "  wire  contains 
these  elements  and  usually  some  iron.  In  order  to  prevent  loss  of 
heat  the  furnace  is  covered  with  asbestos,  which  is  a  poor  con- 
ductor. 

For  technical  use  where  very  high  temperatures  are  required, 
the  resistor  is  usually  powdered  coke,  which  offers  a  great  resistance 
to  the  current.  A  number  of  the  products  manufactured  in  the 
electric  furnace  are  made  from  carbon  and  other  substances,  and 
the  mixture  used  is  itself  the  resistor.  Calcium  carbide  is  made  in 
this  way  from  lime  and  carbon  (217). 

For  high-temperature  work  on  a  small  scale  a  so-called  arc- 
furnace  may  be  used.  In  this  case  the  material  is  placed  in  a  cru- 
cible of  graphite,  which  serves  as  one  electrode;  when  a  carbon  pole 
is  brought  into  contact  with  the  crucible  and  then  separated  from 
it,  an  arc  is  produced  and  the  crucible  is  heated  to  a  very  high  tem- 
perature as  long  as  the  arc  is  maintained. 

So-called  induction  furnaces  are  used  in  the  arts  to  melt  iron 
and  other  metals. 

581.  (b)   Use  of  the  Electric  Current  in  Preparing  Elements  and 
Compounds. — The  more  active  of  the  metals  are  prepared  by  the 
electrolysis  of  their  compounds,  on  account  of  the  fact  that  their 
oxides  are  not  reducible  by  carbon  at  the  temperature  produced 
when  the  latter  unites  with  oxygen.     Sodium,  potassium,  calcium, 
magnesium,    and    aluminium   are    obtained    in    this   way.     The 
details  of  the  methods  used  will  be  described  later.     The  active 
non-metallic  elements,  chlorine  and  fluorine,  are  also  obtained 
by  electrolysis,  as  we  have  seen.     Sodium  hydroxide  and  the 
hydroxides  of  the  other  alkali  metals  are  prepared  industrially 
from  salts  by  electrolytic  methods. 

582.  (c)   Use  of  the  Electric  Current  in  the  Purification  of  Metals. 
—The  presence  of  small  amounts  of  foreign  substances  in  metals 


ELECTROCHEMISTRY  495 

often  affects  materially  their  properties.  The  electrical  conduc- 
tivity of  copper,  for  example,  is  markedly  affected  in  this  way,  and, 
as  a  consequence,  it  is  brought  to  a  high  state  of  purity  before  being 
used  for  electrical  'purposes.  This  is  most  readily  accomplished 
by  depositing  the  metal  from  a  solution  of  its  salt  by  means  of  an 
electric  current.  The  crude  copper  as  it  comes  from  the  smelter 
is  suspended  in  a  bath  containing  a  solution  of  copper  sulphate, 
and  serves  as  one  pole;  the  other  is  a  thin  plate  of  the  pure  metal. 
The  poles  are  connected  with  a  source  of  electricity  in  such  a  way 
that  the  metal  will  be  deposited  on  the  pure  copper.  The  copper 
ions  are  deposited  on  the  negative  pole  and  copper  passes  into 
solution  at  the  positive  pole.  If  the  crude  copper  contains  any 
metal  above  copper  in  the  electromotive  series,  such  as  zinc,  it  will 
pass  into  solution;  elements  less  active  than  copper,  such  as  gold 
and  silver,  will  not  be  converted  into  ions  but  will  fall  to  the  bottom 
of  the  tank  along  with  compounds  of  copper  with  non-metals,  such 
as  the  sulphide  and  arsenide.  At  the  cathode  the  metals  more 
active  than  copper  will  not  be  deposited,  because  the  tendenc}r  of 
copper  ions  to. lose  their  charge  and  pass  into  the  metallic  condi- 
tion is  greater  than  that  of  the  elements  above  it  in  the  electro- 
motive series.  As  far  as  the  copper  is  concerned,  the  reaction  at 
one  pole  is  just  the  reverse  of  that  at  the  other;  the  energy  required 
to  set  copper  free  from  its  ions  at  the  cathode  is  equal  to  that  lib- 
erated by  the  ionization  of  the  metal  at  the  anode.  Consequently, 
the  energy  required  is  only  that  necessary  to  overcome  the  resistance 
of  the  flow  of  the  current  through  the  solution,  which  results  in  the 
movement  of  the  ions.  This  is  reduced  by  keeping  the  liquid  in 
the  cell  well  stirred  and  having  the  plates  placed  close  together. 
In  practice  the  electrolysis  is  carried  out  in  a  trough  lined  with  lead 
into  which  are  placed  alternately  the  anodes  and  cathodes.  Under 
these  conditions  an  electromotive  force  of  less  than  0.5  volt  is 
required.  The  copper  obtained  in  this  way  is  about  99.8  per  cent 
pure.  Very  large  amounts  of  gold  and  silver  are  recovered  from 
the  sludge  obtained  from  the  tank.  The  process  is  also  the  chief 
source  of  tellurium,  which  is  obtained  from  the  copper  telluride 
present  in  the  crude  copper. 

Iron  which  is  to  be  used  in  the  cores  of  electro-magnets  must 
be  as  free  as  possible  from  impurities.     The  metal  can  be  purified 


496  INORGANIC  CHEMISTRY  FOR  COLLEGES 

for  this  purpose  by  a  method  analogous  to  that  used  in  the  case  of 
copper,  and  when  the  highest  purity  is  desired  this  method  is 
employed. 

583.  Electroplating. — This  process  can  be  carried  out  in  a 
manner  analogous  to  that  described  for  the  purification  of  copper. 
The  cathode  is  the  object  to  be  coated,  the  anode  the  metal  to  be 
used,  and  the  electrolyte  a  solution  which  furnishes  the  ions  of 
the  metal  at  such  a  concentration  that  when  they  deposit,  a  closely 
adhering  layer  of  the  metal  is  formed. 

Silver  and  gold  are  commonly  used  in  electroplating.  As 
these  elements  are  very  inactive,  they  are  deposited  from  solutions 
of  most  of  their  salts  when  brought  into  contact  with  more  active 
metals.  If  iron,  for  example,  is  placed  in  a  solution  of  silver 
nitrate,  the  metal  goes  into  solution  and  silver  is  deposited  in  the 
form  of  a  loose  powder  on  the  surface  of  the  metal.  It  is  evident, 
therefore,  that  silver  nitrate  could  not  be  used  as  the  electrolyte  in 
electroplating  silver,  for  the  metal  would  be  deposited  chemically 
in  a  non-adhering  form.  A  salt  must  be  used,  accordingly,  which 
produces  such  a  low  concentration  of  silver  ions,  that  the  ten- 
dency of  the  latter  to  pass  out  of  solution  in  the  presence  of  the 
more  active  metal  is  reduced  to  the  point  where  silver  does  not 
separate;  the  double  cyanide  of  silver  and  potassium,  KAg(CN)2, 
is  such  a  salt.  When  this  compound  is  dissolved  in  water  the  ions 
formed  are  chiefly  K+  and  Ag(CN)2~,  but  a  very  small  concen- 
tration of  silver  ions  is  produced  as  the  result  of  the  further 
ionization  of  the  negative  Ag(CN)2~  ion  into  Ag+  and  2CN~. 
The  concentration  of  the  silver  ion  is  so  small  that  its  direct 
deposition  on  the  metal  without  the  use  of  the  current  does  not 
take  place.  Even  under  these  conditions  it  is  difficult  to  deposit 
silver  electrolytically  on  iron  and  the  metal  is  usually  first  coated 
with  copper  upon  which  the  silver  is  deposited.  Double  cyanides 
are  also  used  in  electroplating  with  gold  and  with  nickel. 

In  making  electrotypes  for  use  in  a  printing  press,  a  wax  mold  is 
first  made  with  the  impression  desired.  Its  surface  is  next  coated 
with  finely  powdered  graphite  to  make  it  a  conductor  of  electricity. 
It  is  then  suspended  in  a  bath  of  a  copper  salt  and  connected  with 
the  negative  pole  of  a  supply  of  electricity.  A  plate  of  copper  in 
the  bath  is  connected  with  the  positive  pole.  When  the  deposit 
of  copper  on  the  mold  is  sufficiently  thick,  the  metal  is  stripped 


ELECTRO-CHEMISTRY  497 

from  the  wax,  and  backed  with  stereotype    metal,    and   finally 
mounted  on  a  wooden  block. 

584.  Oxidation  and  reduction  are  carried  out  by  means  of  the 
electric  current,  because  it  is  possible  to  obtain  hydrogen  and 
oxygen  when  aqueous  solutions  are  electrolyzed.  It  has  been 
stated,  for  example,  that  perchloric  acid,  HCICU,  is  now  made 
by  the  electrolytic  oxidation  of  hydrochloric  acid,  and  potassium 
chlorate  in  an  analogous  way  from  potassium  chloride.  Elec- 
trolytic oxidation  and  reduction  are  conveniently  applied  in  the 
manufacture  of  certain  important  organic  compounds. 

EXERCISES 

1.  The    reactions    Zn^Zn++-f20     and    H2  ^±  2H+  -f  2  ©     are 
reversible.     Show  how  the  facts  in  regard  to  these  changes  are  in  accord 
with  the  law  of  mobile  equilibrium. 

2.  Draw  a  diagram  indicating  the  flow  of  the  electric  current  and  the 
changes  which  take  place  when  zinc  is  in  contact  with  copper  in  the  presence 
of  a  solution  of  hydrochloric  acid. 

3.  Draw  a  diagram  to  illustrate  how  an  experiment  could  be  carried  out 
to  test  Faraday's  law. 

4.  What  weights  of  copper  and  chlorine  will  be  liberated  when   (a)  20,000 
coulombs  are  passed  through  a  solution  of  copper  chloride,  and  (6)  when  a 
current  of  5  arnperes  is  passed  for  three  hours  through  the  solution. 

5.  An  electric  current  was  passed  first  through  a  solution  of  a  silver  salt 
and  then  through  one  of  a  copper  salt  in  which  the  metal  had  the  valence  2. 
It  was  found  that  in  the  same  time  2.3733  grams  Ag  and  0.6992  gram  Cu  were 
deposited.     Assuming  107.88  as  the  atomic  weight  of  silver  calculate  the 
atomic  weight  of  copper. 

6.  Complete  the  following  expressions,  replacing  a,  6,  c,  etc.,  by  the  proper 
units:    (a)    number  of  coulombs  =  aX&;    (6)    number  of  joules  —  cXd;    (c) 
number  of  kilowatt-hours  =  eXf=gXhXi. 

7.  It  requires  2.3  volts  to  liberate  chlorine  from  a  solution  of  sodium 
chloride,     (a)  Calculate  the  number  of  joules  required  to  set  free  1  gram- 
atomic-weight  of  chlorine:    (6)  how  many  kilowatt-hours  is  this  number  of 
joules?     (c)  If  electrical  energy  costs  10  cents  per  kilowatt-hour,  what  is  the 
cost  of  the  electrical  energy  required  to  set  free  1  kilo  of  chlorine? 

8.  What  does  the  electrical  energy  cost  to  run  a  60-candle-power  tungsten 
incandescent  lamp  for  1  hour  if  the  energy  cost  10  cents  per  kilowatt-hour 
and  the  lamp  consumes  1 . 2  watts  per  candle  power? 

9.  Calculate  the  E.M.F.  of  cells  made  up  of  the  following  pah's  of  elements 
in  contact  with  normal  solutions  of  the  respective  salts,  and  state  in  which 
direction  the  positive  current  would  flow  in  each  case  through  a  wire  connect- 
ing the  two  metallic  poles:   (a)  Al  and  Zn,    (6)  Zn  and  Co,    (c)  Zn  and  Pb, 
(d)  ZnandAu,    (e)  Hg  and  Ag. 


498  INORGANIC  CHEMISTRY  FOR  COLLEGES 

10.  Explain  how  you  can  tell  by  the  determination  of  the  specific  gravity 
of  the  sulphuric  acid  in  a  lead  storage  cell  whether  or  not  it  is  charged. 

11.  The  resistance  per  foot  of  number  20  nichrome  wire,  which  is  0.81  mm. 
in  diameter,  is  0.52  ohm.     (a)  How  many  feet  of  the  wire  are  required  to 
furnish  a  resistance  of  100  ohms?     (6)  If  the  difference  of  potential  at  the 
ends  of  a  wire  having  a  resistance  of  100  ohms  is  110  volts  how  many  cou- 
lombs pass  through  the  wire  per  second  (current  =  E.M.F.  -r-  resistance)? 
(c)  How  many  calories  will  be  obtained  if  the  current  passes  for  one  hour? 

12.  Calculate  the  cost  of  1,000,000  calories  produced  by  (a)  burning  coal 
at  $10  per  ton  (1  gram  coal  yields  8000  calories)  and  by  (6)  the    electric 
current  at  2  cents  per  kilowatt-hour. 

13.  Two  dry  cells  were  tested  by  applying  the  terminals  of  a  voltmeter 
to  the  zinc  and  carbon  poles  of  the  cell.     In  each  case  the  meter  registered 
1.9  volts.     When  tested  with  an  ammeter  each  cell  registered  30  amperes, 
(a)  If  the  carbon  poles  of  the  two  cells  are  connected  to  each  other  and  the 
zinc  poles  are  connected,  what  would  the  voltmeter  and  the  ammeter  register 
if  joined  to  the  carbon  and  the  zinc  poles?     (6)  If  the  carbon  pole  of  one  cell 
is  connected  with  the  zinc  pole  of  the  other  and  the  measuring  instruments 
connected  with  the  free  poles  what  would  the  readings  be  in  each  case?     (c) 
If  six  partially  used  cells  were  available  which,  when  tested,  showed  1.6  volts 
and  8  amperes  how  should  they  be  connected  to  use  with  a  small    incan- 
descent lamp  requiring  3  volts? 

14.  When  the  salt  of  the  formula  KAg(CN)2  is  used  in  electroplating  why 
is  potassium  not  deposited  at  the  cathode? 

15.  How  could  gold  be  separated  from  zinc  by  electrolysis? 


CHAPTER  XXXIV 
THE  PROPERTIES  OF  OXIDES,  HYDROXIDES,  AND  SALTS 

585.  A  very  large  number  of  compounds  are  known  which 
contain  metallic  atoms.     With  the  aid  of  the  electromotive  series 
and  the  periodic  law,  however,  it  is  possible  to  systematize  the 
facts  to  be  presented,  by  grouping  the  compounds  into  several 
classes,  the  members  of  which  exhibit    similar  properties  that 
vary  progressively  with  the  atomic  weight  and  activity  of  the 
metal.     In  this  chapter  the  properties  of  compounds  and  their 
relationships  will  be  considered  in  a  general  way,  and  later  specific 
facts  will  be  given  in  regard  to  those  particular  compounds  which 
are  of  interest   from  a  theoretical   point  of  view  or  have  found 
important  practical  applications. 

OXIDES 

586.  All  the  metals  form  well  characterized  oxides,  and  in  the 
case  of  most  metals  two  or  more  compounds  of  this  class  are  kriown. 
The  highest  valence  which  can  be  shown  by  a  metal  toward  oxygen 
is  indicated  by  its  position  in  the  periodic  classification,  and  in 
most  cases  oxides  are  known  in  which  this  valence  is  shown;  the 
most  striking  exceptions  occur  in  the  case  of  the  metals  in  the 
eighth  group,  where  osmium  and  ruthenium  alone  form  oxides  in 
which  the  metals  have  the  valence  8. 

Oxides  are  formed  by  heating  metals  in  the  air  or  oxygen,  or 
by  decomposing  salts  of  the  metals.  Whether  or  not  a  metal 
burns  in  the  air  is  determined  in  certain  cases  by  its  physical  con- 
dition. Under  ordinary  circumstances  the  metals  in  the  elec- 
tromotive series  down  to  and  including  zinc  burn  in  the  air;  iron 
and  lead  can  be  obtained  in  such  a  finely  divided  condition  that 
they  take  fire  spontaneously  when  brought  into  the  air.  The 
metals  down  to  silver  will  burn  in  oxygen.  The  composition  of 

499 


500  INORGANIC  CHEMISTRY  FOR  COLLEGES 

the  oxide  formed  when  a  metal  burns  is  determined  by  the  activity 
of  the  latter  and  its  valence.  Sodium  and  potassium  form  per- 
oxides, Na202  and  K^O*,  which  contain  an  excess  of  oxygen  over 
that  present  in  the  normal  oxides,  Na2O  and  K2O.  The  alkaline 
earths,  magnesium,  zinc,  cobalt,  nickel,  tin,  lead,  and  copper  give 
oxides  of  the  general  formula  MO;  aluminium,  chromium,  arsenic, 
antimony,  and  bismuth,  oxides  of  the  formula  M^Os;  and  man- 
ganese and  iron,  oxides  of  the  formula  MsCX  in  which  two  atoms 
of  the  metal  have  the  valence  3  and  one  the  valence  2. 

587.  The  chemical  properties  of  an  oxide  of  a  metal  are 
determined  by  the  activity  of  the  metal  and  its  valence.  Some 
oxides  react  with  acids  to  form  salts,  some  with  bases  to  form 
salts,  and  some  in  both  ways;  other  oxides  react  with  neither 
acids  nor  bases.  The  oxides  of  the  more  active  metals — the  alkali 
metals  and  those  of  the  alkaline  earths — dissolve  in  water,  and 
form  strong  bases,  which  react  with  acids  to  form  salts.  Mag- 
nesium oxide  is  also  basic.  Aluminium  oxide  dissolves  in  both 
acids  and  bases,  and,  therefore,  exhibits  both  acidic  and  basic 
properties.  Certain  metals  which  can  show  a  high  valence  form 
oxides  in  which  they  function  as  acid-forming  elements.  Man- 
ganese, for  example,  which  is  in  the  seventh  group  in  the  periodic 
classification,  forms  an  oxide  of  the  formula  M^Oy,  which  is  the 
anhydride  of  permanganic  acid,  HMnO-i;  the  latter  is  a  strong 
acid  and  forms  salts  which  resemble  in  composition  and  properties 
those  derived  from  perchloric  acid,  HCICU.  As  has  already  been 
emphasized,  increase  in  valence  of  an  element  toward  oxygen  is 
associated  with  the  development  of  acidic  properties;  and  this 
occurs  even  when  the  element  is  a  metal.  Manganese  is  like  iron 
in  many  of  its  physical  and  chemical  properties.  When  it  has  the 
valence  2  it  acts  as  a  strong  base-forming  element,  with  the  valence 
3  it  is  still  base-forming  but  less  active. 

The  temperature  at  which  the  oxides  are  reduced  by  hydrogen 
or  carbon  varies  with  the  activity  of  the  metal  and  the  valence  it 
shows  in  the  oxide.  When  heated  with  hydrogen  the  oxides 
down  to  and  including  manganese  can  be  reduced  to  lower  oxides, 
if  they  exist,  but  complete  reduction  to  the  metal  cannot  be 
effected.  The  oxides  of  cadmium  and  the  metals  below  it  can 
be  easily  reduced  to  the  metal  by  hydrogen.  The  oxides  of  mer- 
cury and  the  less  active  metals  are  readily  decomposed  by  heat 


THE  PROPERTIES  OF  OXIDES,  HYDROXIDES,  AND  SALTS     501 

alone.  At  the  temperature  obtainable  in  a  blast-furnace,  the 
oxides  of  the  metals  from  manganese  down  can  be  reduced  by 
carbon. 

HYDROXIDES 

588.  The  hydroxides  of  the  alkali  metals  are  soluble  in  water 
and  are  formed  by  the  action  of  water  on  the  oxides  of  the  metals; 
they  are   prepared   by  the   electrolysis   of   the   chlorides.     The 
hydroxides  of  the  alkaline  earths  are  difficultly  soluble  in  water 
and  are  made  by  the  action  of  water  on  the  oxides.     The  hydrox- 
ides of  the  other  metals,  being  insoluble,  are  formed  as  precipi- 
tates when  the  salts  of  the  metals  are  treated  with  a  solution  of 
a  soluble  hydroxide.    When,  however,  a  soluble  base  is  added  to  a 
solution  of  a  salt  of  silver,  mercury,  cuprous  copper  Cu1,  or  aurous 
gold  Au1,  an  insoluble  oxide  is  formed.     In  the  case  of  silver, 
some  of  the  oxide  dissolves  as  hydroxide. 

All  the  hydroxides  except  those  of  the  alkali  metals  are  decom- 
posed into  oxides  and  water  when  heated.  They  all  dissolve  in 
acids  to  form  salts.  The  hydroxides  of  univalent  metals  are 
strong  bases,  and,  as  a  result,  form  neutral  salts  with  strong 
acids  that  are  not  hydrolyzed  by  water.  Cuprous  salts  and  aurous 
salts  are  hydrolyzed.  The  hydroxides  of  the  bivalent  metals  are 
relatively  strong  bases  and  their  salts  are  hydrolyzed  but  slightly 
if  at  all  at  ordinary  temperatures.  The  hydroxides  of  the  tri- 
valent  metals  are  weak  bases;  their  salts  with  strong  acids  are 
appreciably  hydrolyzed  and  those  with  weak  acids,  such  as  car- 
bonic acid  and  hydrogen  sulphide,  are  completely  hydrolyzed  by 
water. 

589.  Certain  hydroxides  which  are  precipitated  by  a  solution 
of  sodium  hydroxide  are  not  precipitated  by  one  of  ammonium 
hydroxide  if  ammonium  salts  are  present.     This  fact  is  utilized 
in  analytical  chemistry  in  quantitative  separations  of  metals  and 
is,  therefore,  of  importance.     The  explanation  is  based  on  a  con- 
sideration of  the  facts  from  the  point  of  view  of  the  law  of  molecu- 
lar concentration. 

When  a  hydroxide  is  precipitated  from  solution,  the  com- 
pound is  formed  as  the  result  of  the  union  of  the  metallic  and  the 
hydroxyl  ions.  For  example,  in  the  case  of  a  bivalent  metal  the 
reaction  may  be  indicated  as  follows:  M+  ++20H~^±M(OH)2. 


502  INORGANIC  CHEMISTRY  FOR  COLLEGES 

If  the  solubility  of  the  hydroxide  is  very  small,  the  amount  in 
solution  is  completely  ionized  and  an  equilibrium  is  set  up  between 
the  solid  present  and  its  ions.  According  to  the  law  of  molecular 
concentration,  in  the  case  of  the  hydroxide  of  a  bivalent  metal 
the  concentration  of  the  metallic  ions  multiplied  by  the  square  of 
the  concentration  of  the  hydroxyl  ions  is  a  constant.  Similar 
equilibria  exist  between  the  ions  of  other  slightly  soluble  com- 
pounds; the  constant  for  such  an  equilibrium  is  commonly  called 
the  solubility  product.  If  the  ions  of  a  slightly  soluble  compound 
are  brought  together  in  solution,  they  will  react  to  form  the  com- 
pound, which  will  precipitate  until  the  solubility  product  of  the 
ions  is  reached.  For  example,  when  sodium  hydroxide,  which 
furnishes  a  high  concentration  of  OH~  ions;  is  added  to  a  strong 
solution  of  calcium  chloride,  calcium  hydroxide  is  precipitated 
because  the  solution  furnishes  a  higher  concentration  of  OH"  and 
Ca++  ions  than  that  indicated  by  the  solubility  product  of  these 
ions  in  equilibrium  with  calcium  hydroxide.  When  ammonium 
hydroxide,  which  is  a  weak  base,  is  added  to  a  solution  of  a  cal- 
cium salt,  the  solubility  product  is  not  exceeded  and  the  calcium 
hydroxide  is  not  precipitated.  Ammonium  hydroxide  does  pre- 
cipitate magnesium  hydroxide,  however;  the  solubility  of  the 
latter  is  much  less  than  that  of  calcium  hydroxide,  and,  as  a  result, 
the  solubility  product  of  magnesium  and  hydroxyl  ions  is  less. 
When  ammonium  hydroxide  is  added  to  a  solution  of  a  mag- 
nesium salt,  magnesium  hydroxide  is  precipitated  until  the  concen- 
trations of  the  OH~  and  Mg++  ions  are  reduced  to  those  indicated 
by  the  solubility  product.  If  a  solution  containing  ammonium 
hydroxide  and  ammonium  chloride  is  added  to  a  magnesium  salt, 
the  hydroxide  is  not  precipitated.  This  is  due  to  the  fact  that  the 
extent  of  ionization  of  ammonium  hydroxide,  and,  therefore,  the 
concentration  of  hydroxyl  ions  in  its  solution,  is  reduced  in  the 
presence  of  ammonium  salts.  There  is  an  equilibrium  established 
when  ammonium  hydroxide  ionizes: 


If  ammonium  ions,  produced  from  ammonium  chloride, 
NH4C1  = 


THE  PROPERTIES  OF  OXIDES,  HYDROXIDES,  AND  SALTS     503 

are  added  to  the  solution  of  ammonium  hydroxide,  the  equilibrium 
in  the  latter  case  is  disturbed  and  shifts  from  right  to  left;  as  a 
result,  NH4+  and  OH~  unite  to  form  NH4OH,  and  the  concen- 
tration of  the  OH~  ions  is  reduced.  If  a  sufficient  quantity 
of  an  ammonium  salt  is  added  to  a  solution  of  ammonium  hydrox- 
ide, the  concentration  of  the  OH~  ions  is  reduced  to  such  an  extent 
that  in  the  presence  of  a  magnesium  salt  the  solubility  product 
of  magnesium  and  hydroxyl  ions  is  not  obtained  and  the  precipi- 
tation of  magnesium  hydroxide  does  not  take  place. 

The  precipitation  of  ferric  hydroxide,  Fe(OH)a,  by  ammonium 
hydroxide  is  not  prevented  by  the  presence  of  ammonium  salts 
because  the  solubility  of  the  metallic  hydroxide  is  very  small, 
and  it  furnishes,  therefore,  a  very  low  concentration  of  hydroxyl 
ions.  It  is  impossible  by  the  addition  of  ammonium  salts  to 
reduce  the  concentration  of  hydroxyl  ions  in  a  solution  of 
ammonium  hydroxide  to  such  an  extent  that  the  solubility  prod- 
uct of  ferric  and  hydroxyl  ions  is  not  exceeded;  precipitation  of 
ferric  hydroxide,  accordingly,  occurs.  Ferrous  hydroxide,  on  the 
other  hand,  is  more  soluble,  and  its  precipitation  by  ammonium 
hydroxide  is  prevented  by  ammonium  salts.  The  solubilities  of 
the  hydroxides  of  most  of  the  metals  having  the  valence  2  are 
such  that  they  are  not  precipitated  by  ammonium  hydroxide  in 
the  presence  of  ammonium  salts;  the  hydroxides  of  the  trivalent 
metals  are  much  less  soluble  and,  as  a  consequence,  they  are 
precipitated  under  these  conditions. 

SALTS 

•  590.  The  methods  of  preparing  salts  (142, 148)  and  some  of  the 
properties  which  they  possess  in  general  (241,  242,  246,  249,  251) 
have  already  been  described  at  some  length.  The  solubilities 
of  the  salts  of  the  more  important  acids  which  have  been  con- 
sidered, have  been  stated  and  the  action  of  heat  on  acid,  basic, 
and  neutral  salts  has  been  emphasized  (204,  282,  294,  307,  367, 
373,  408,  428,  460,  462,  511). 

591.  Isomorphism. — When  a  salt  crystallizes,  the  form  it 
assumes  is  characteristic  of  it.  While  the  size  and  general  shape 
of  the  crystal  may  vary  according  to  the  conditions  under  which 
crystallization  takes  place,  the  angles  at  which  the  faces  cut  one 


504  INORGANIC  CHEMISTRY  FOR  COLLEGES 

another  are  fixed.  In  general,  on  the  slow  evaporation  of  a  solu- 
tion containing  two  or  more  salts  the  latter  will  be  deposited  as 
individual  crystals.  In  the  case  of  certain  salts,  however,  which 
are  closely  alike  in  crystalline  form,  crystals  of  but  one  kind  will 
be  produced,  which  contain  the  salts  in  the  proportion  in  which 
they  were  present  in  the  solution.  Salts  which  produce  such  mixed 
crystals  are  said  to  be  isomorphous.  In  general,  isomorphous 
salts  are  closely  related  in  composition;  for  example,  zinc  sul- 
phate, ZnSC>4,7H2O,  and  manganese  sulphate,  MnSO4,7H2O, 
are  isomorphous,  and  it  is,  therefore,  impossible  to  separate  them 
by  crystallization. 

592.  Effect  of  Temperature  on  the  Solubility  of  Salts.— The 
change  in  the  solubility  of  salts  with  change  in  temperature  is  a 
fact  of  great  importance,  which  is  utilized  in  separating  salts  and 
in  their  purification.  Crytallization  is  one  of  the  most  effective 
processes  used  in  the  preparation  of  soluble  chemical  compounds. 

The  change  in  solubility  with  rise  in  temperature  is  markedly 
different  in  the  case  of  different  salts;  for  example,  under  these 
conditions  the  solubility  of  potassium  nitrate  increases  greatly, 
that  of  sodium  chloride  only  slightly,  whereas  anhydrous  sodium 
sulphate  is  more  soluble  in  cold  than  in  hot  water.  These  differ- 
ences in  the  effect  of  increased  temperature  on  the  solubility  of 
salts  can  be  traced  to  the  energy  changes  which  take  place  when 
solution  occurs.  When  potassium  nitrate  is  dissolved  in  water  the 
temperature  of  the  solution  falls  far  below  that  of  the  solvent; 
much  heat  is  absorbed  from  the  surroundings: 

Solid  KNO3  +  water  =  solution  of  KNO3  -  8500  cal. 

This  equation,  which  expresses  the  heat  of  solution  of  potassium 
nitrate,  may  be  abbreviated  to  read,  KNOs+Aq  =  —8500  cal. 
It  means  that  when  1  gram-molecular-weight  of  the  salt  is 
dissolved  in  such  an  amount  of  water  (aqua)  that  further  addition 
of  the  solvent  produces  no  heat  change,  the  heat  absorbed  is  8500 
calories. 

There  is  an  equilibrium  between  the  solid  and  the  solution,  and 
since  the  transfer  of  the  salt  in  either  direction  is  accompanied  by 
a  heat  change,  the  law  of  van't  Hoff  must  apply.  It  will  be 
recalled  that  rise  in  temperature  shifts  the  equilibrium  so  that 


THE  PROPERTIES  OF  OXIDES,  HYDROXIDES,  AND  SALTS     505 

heat  is  absorbed;    as  a  consequence,  under  these  conditions  the 
salt  becomes  more  soluble. 

In  the  case  of  sodium  chloride  the  change  is  in  the  same  direc- 
tion as  with  potassium  nitrate,  for  NaCl+Aq  =  -1200  cal;  but 
the  heat  change  is  much  smaller  and  the  effect  of  rising  tempera- 
ture is,  as  a  result,  less.  Anhydrous  sodium  sulphate  dissolves  in 
water  with  rise  in  temperature:  NaoSCU+Aq  =  +460  cal.  In 
accordance  with  the  law  of  mobile  equilibrium  the  solubility  of 
the  salt  decreases  with  rise  in  temperature.  Other  changes  than 
solution  occur  when  certain  salts  dissolve  in  water.  Among  these 
are  hydration,  hydrolysis,  and  ionization;  and  as  these  processes 
produce  heat  effects,  the  change  in  solubility  of  a  salt  with  change 
in  temperature  cannot  always  be  determined  by  a  consideration  of 
its  heat  of  solution.  In  the  diagram  on  page  506  are  plotted  the 
solubility  curves  of  a  few  salts.  The  concentrations  are  expressed 
as  the  number  of  grams  of  the  anhydrous  salt  which  are  held  in 
solution  by  100  grams  of  water.  The  slope  of  the  curve  in  each  case 
indicates  the  change  in  solubility  with  rise  in  temperature;  that 
of  sodium  chloride  is  nearly  horizontal,  whereas  the  curve  of 
potassium  rises  steeply. 

Advantage  is  taken  of  the  fact  that  salts  absorb  heat  when  they 
pass  into  solution,  in  making  freezing  mixtures  for  low  tempera- 
ture work.  A  few  examples  of  such  mixtures  in  which  the  ingre- 
dients are  expressed  as  parts  by  weight  are  as  follows:  30  NEUCl 
and  100  water  at  13.3°,  temperature  reduced  to  -5.1°;  60 
NH4N03  and  100  water  at  13.6°,  temperature  reduced  to  -13.6°; 
110  Na2S2O3,5H2O  and  100  water  at  10.7°,  temperature  reduced 
to  -8°;  250  CaCl2,6H2O  and  100  water  at  10.8°,  temperature 
reduced  to  -12.4°. 

When  the  effect  of  increased  temperature  on  the  solubility  of 
two  salts  is  markedly  different,  it  is  easy  to  separate  them  from 
each  other  by  crystallization.  A  case  of  this  kind  is  potassium 
nitrate  and  sodium  chloride.  If  water  is  saturated  with  the  two 
salts,  say  at  60°,  and  then  cooled,  to  say  10°,  but  a  small  quantity  ' 
of  sodium  chloride  will  crystallize  out,  since  its  solubility  at  60° 
is  about  38  grams  in  100  of  water  and  about  36  grams  at  10°. 
In  the  case  of  potassium  nitrate  the  solubilities  at  these  tem- 
peratures are  110  and  17  grams,  respectively,  and  a  large  part 
of  the  salt  will  separate  when  the  solution  cools.  If  the  mixture 


506 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


of  the  two  salts  obtained  is  recrystallized  from  hot  water,  the 
potassium  nitrate  which  separates  is  practically  free  from  salt. 

It  is  seen  from  the  diagram  that  at  room-temperatures  potas- 
sium nitrate  is  much  less  soluble  than  sodium  nitrate.     Conse- 


30        40         50         60        70 
Temperature  ,Deg.  Cent. 

FIG.  36. 

quently,  if  boiling  water  is  saturated  with  sodium  nitrate  and 
potassium  chloride  and  the  solution  allowed  to  cool,  a  double 
decomposition  takes  place  and  potassium  nitrate  crystallizes  out. 
This  fact  is  utilized  in  manufacturing  potassium  nitrate  from  Chile 
saltpeter.  In  the  great  salt  beds  at  Stassfurt  and  other  places  the 


THE  PROPERTIES  OF  OXIDES,  HYDROXIDES,  AND  SALTS     507 

salts  are  found  in  the  mines  deposited  in  layers,  the  positions  of 
which  are  in  accord  with  the  solubilities  of  the  several  salts. 
This  fact  shows  that  they  were  deposited  as  the  result  of  the 
evaporation  of  an  aqueous  solution. 

593.  The  Solubility  of  Salts  in  Acids.—  The  difficultly  soluble 
salts  of  weak  acids  are  more  or  less  soluble  in  strong  acids,  for 
example,  calcium  phosphate,  which  is  practically  insoluble  in 
water,  dissolves  in  nitric  acid.  The  explanation  of  this  fact  is 
based  on  a  consideration  of  the  reactions  which  take  place  between 
the  ions  present  in  the  solution.  When  calcium  phosphate  is  in 
contact  with  water  a  very  small  amount  of  the  salt  passes  into 
solution  and  forms  ions: 


+  2PO4  --- 
Nitric  acid  furnishes  ions  also  : 

HN03  ^  H+  +  N03" 

When  phosphoric  acid  is  in  solution  it  ionizes  almost  wholly  as 
follows  : 

^H+  -f  H2P04~ 


Just  as  hydrogen  ions  and  hydroxyl  ions  when  brought  together 
unite  to  form  water  —  H+  +  OH~  +±  HoO  —  so  hydrogen  ions 
unite  with  P04~"  ~  ions  to  form  the  ion 


This  occurs  when  calcium  phosphate  is  placed  in  nitric  acid;  the 
removal  of  the  PO4~  "  ions  disturbs  the  equilibrium  between 
the  solid  calcium  phosphate  and  its  solution,  and  more  of  the  salt 
passes  into  solution  until  it  finally  dissolves.  The  ions  finally 
present  are  Ca++;  NO3~,  H2PO4~,  and  H+. 

In  general,  a  salt  of  a  weak  acid  will  dissolve  in  a  strong  acid 
as  the  result  of  the  fact  that  the  hydrogen  ions  furnished  by  the 
strong  acid  unite  with  the  radical  of  the  weak  acid  to  form  the 
undissociated  free  acid  or,  as  in  the  case  of  phosphoric  acid,  an 
undissociated  ion.  The  cause  of  the  change  is  the  same  as  that 
which  leads  to  the  neutralization  of  a  base  by  an  acid.  If  in  any 
way  the  ions  of  a  very  slightly  dissociated  substance  are  brought 
together,  they  will  unite  until  their  concentrations  are  those  which 


508  INORGANIC  CHEMISTRY  FOR  COLLEGES 

can  exist  in  equilibrium  with  the  undissociated  substance,  that  is, 
until  the  solubility  product  is  reached. 

Whether  or  not  the  salt  of  a  weak  acid  will  dissolve  in  a  strong 
acid  is  determined  not  only  by  the  relative  strength  of  the  acids 
concerned,  but  also  by  the  solubility  of  the  salt  in  water.  This 
can  be  shown  by  a  consideration  of  the  sulphides  of  zinc  and 
copper.  In  the  case  of  the  former  a  certain  amount  of  the  salt 
passes  into  solution  and  ionizes: 

ZnS^Zn+  +  +  S- 

When  hydrochloric  acid  is  added,  hydrogen  ions  are  present  and 
an  equilibrium  is  set  up  between  them  and  the  sulphur  ions: 

2H+  +  S" 

If  the  concentration  of  the  S~~  ions  furnished  by  the  zinc  sul- 
phide is  greater  than  that  of  these  ions  when  in  equilibrium  with 
H^~  ions,  the  equilibrium  will  shift  and  H^S  molecules  will  be 
formed.  More  zinc  sulphide  will  pass  into  solution  as  the  result 
of  the  removal  of  S —  ions,  and  finally  the  salt  will  dissolve.  In 
the  case  of  copper,  the  sulphide  is  so  little  soluble  that  the  S~~ 
ions  are  not  removed,  and  the  sulphide  does  not  dissolve.  It  is 
for  the  reasons  just  stated  that  hydrogen  sulphide  will  precipitate 
copper  sulphide  from  a  solution  of  one  of  its  salts  in  the  presence 
of  hydrochloric  acid,  but  will  not  precipitate  zinc  sulphide  under 
the  same  conditions.  Zinc  sulphide  is  precipitated,  however,  by 
ammonium  sulphide,  because  the  solution  contains  no  hydrogen 
ions. 

594.  The  Use  of  Sulphides  in  Qualitative  Analysis. — The 
difference  in  the  solubility  of  the  sulphides  in  water,  and  in  acids, 
is  utilized  in  qualitative  analysis.  The  sulphides  of  the  alkali 
metals  are  soluble  in  water,  and  those  of  the  alkaline  earths  and 
magnesium  are  hydrolyzed  by  water;  they  are  not,  as  a  conse- 
quence, precipitated  by  hydrogen  sulphide.  The  sulphides  of  the 
metals  in  the  electromotive  series  from  magnesium  to  iron,  inclu- 
sive, when  they  are  in  the  bivalent  condition,  are  soluble  in  dilute 
acids;  they  are  precipitated  by  ammonium  sulphide  but  not  by 
hydrogen  sulphide  if  an  acid  is  present.  The  sulphides  of  the 


THE  PROPERTIES  OF  OXIDES,  HYDROXIDES,  AND  SALTS     509 

metals  from  cobalt  to  gold  are  insoluble  in  dilute  acids  and  are 
precipitated  by  hydrogen  sulphide  in  the  presence  of  acids.  It  is 
possible,  therefore,  to  separate  the  metals  into  three  large  groups 
which  can  be  divided  into  subgroups  (Appendix  VI),  and  the 
process  can  be  continued  until  a  complete  separation  has  been 
effected.  The  method  of  doing  this  is  studied  in  detail  in  quali- 
tative analysis. 

595.  The  Solubility  of  Salts  in  a  Solution  of  Ammonia. — Com- 
pounds of  certain  metals  which  are  insoluble  in  water  dissolve  in  a 
solution  of  ammonia.     Among  the  more  important  of  these  are  the 
compounds  of  silver,  zinc,  copper,  nickel,  and  cobalt.     In  all  cases 
ammonia  unites  with  the  metallic  atom  and  substances  more  or 
less  soluble  in  water  are  formed.     Silver  chloride,  for  example,  is 
converted   into   a   compound    of   the   composition   AgfNHs^Cl, 
which  gives  the  ions  Ag(NH3)2+  and  Cl~,  and  is  readily  soluble  in 
water.     In  the  case  of  zinc  4  molecules  of  ammonia  are  added  to 
the  metallic  ion.     Zinc  hydroxide,  which  is  insoluble  in  water,  dis- 
solves in  ammonia  as  the  result  of  the  formation  of  the  compound 
Zn(NH3)4(OH)2.     Cuprous    copper    gives   the   ion   Cu(NH3)2+, 
and  cupric  copper  the  ion  Cu(NH3)4+4~. 

596.  The  Effect  on  the  Solubility  of  a  Salt  of  a  Substance 
Yielding  an  Ion  of  the  Salt. — When  a  salt  is  difficultly  soluble  in 
water  there  is  an  equilibrium  set  up  between  the  solid   and  the 
saturated  solution.     For  example,  in  the  case  of  lead  iodide  the 
equilibrium  can  be  represented  thus, 

PbI2  (solid)  —  Pb++  +2I-, 

since  all  of  the  salt  in  solution  is  ionized.  It  is  evident,  from  an 
application  of  the  law  of  molecular  concentration,  that  if  either 
lead  ions  or  iodine  ions  are  added  to  the  solution  the  equilibrium 
must  shift,  and,  as  a  result,  the  change  which  takes  place  is  that 
indicated  by  reading  the  equation  from  right  to  left;  lead  iodide 
comes  out  of  solution  and  the  solubility  of  the  salt  is  thereby 
reduced. 

The  fact  that  the  solubility  of  salts  is  decreased  in  this  way  is 
utilized  in  analytical  chemistry  because  it  is  advisable  to  reduce  as 
far  as  possible  the  solubility  of  difficultly  soluble  salts  and  thus 
effect  quantitative  separations.  For  example,  in  precipitating 


510  INORGANIC  CHEMISTRY  FOR  COLLEGES 

silver  chloride  in  the  quantitative  determination  of  chlorine  ions,  a 
slight  excess  of  the  silver  salt  is  used  in  the  precipitation  to  reduce 
the  solubility  of  the  silver  chloride  formed. 

597.  The  Solubility  of  Salts  in  Solutions  of  Other  Salts.— In 
certain  cases  salts  which  are  insoluble  in  water  are  soluble  in  solu- 
tions of  other  salts.     Silver  cyanide,  AgCN,  for  example,  is  insol- 
uble in  water  but  dissolves  in  a  solution  of  potassium  cyanide,  KCN, 
as  the  result  of  the  formation  of  a  compound  of  the  composition 
KCN,AgCN.     The  formula  is  preferably  written  in  the  form 
KAg(CN)2  because,  when  it  dissolves  in  water,  the  ions  produced 
are  K~*~  and  Ag(CN)2~.     There  are  many  salts  of  this  type  which 
are  formed  as  the  result  of  the  combination  of  two  salts. 

598.  Double  Salts. — Many  compounds  consisting  of  two  salts 
in  combination  are  formed  from  soluble  salts;  if  potassium  sulphate 
and  aluminium  sulphate  are  dissolved  in  water  and  the  mixture 
of  the  two  solutions  is  allowed  to  evaporate,  crystals  of  a  double 
salt  of  the  composition  K2SO4,Al2 (864)3, 24H2O  are  formed;   the 
compound  is  commonly  called  alum. 

Many  double  chlorides  are  known.  The  one  of  the  composi- 
tion K^PtCle  is  formed  when  potassium  chloride  is  added  to  a  solu- 
tion of  platinic  chloride,  PtCU.  Since  it  is  but  slightly  soluble  in 
water  it  is  used  in  the  quantitative  determination  of  potassium  or 
of  platinum.  The  compound  is  called  potassium  chloroplatinate, 
because  it  may  be  considered  as  derived  from  platinic  acid,  H^PtOs, 
by  the  replacement  of  3  bivalent  oxygen  atoms  by  6  univalent 
chlorine  atoms.  Chloroplatinic  acid,  H^PtClojOH^O,  is  formed 
from  hydrochloric  acid  and  platinic  chloride. 

We  have  already  seen  that  certain  sulphides  unite  to  form  com- 
plex salts  (478,  489)  and  that  they  bear  a  similar  relation  in  com- 
position to  the  salts  of  oxygen  acids;  ammonium  thioarsenate  and 
ammonium  arsenate  have  the  composition  represented  by  the 
formulas  (NH4)aAsS4  and  QTH^sAsQ*.  Many  complex  salts  of 
this  type  are  known.  Since  in  many  of  these  one  metal  functions 
as  an  acid-forming  element,  it  is  clear  why  in  the  majority  of  cases 
complex  salts  are  produced  from  a  salt  of  a  metal  which  can  form 
acids — a  weakly  metallic  element — and  a  salt  of  a  metal  which 
forms  strong  bases. 

The  stability  of  the  double  salts  when  in  contact  with 
water  varies  markedly.  Some,  like  potassium  ferrocyanide, 


THE  PROPERTIES  OF  OXIDES,  HYDROXIDES,  AND  SALTS     511 

Fe(CN)2,4KCN  or  K4Fe(CN)o,  are  stable  and  yield  complex  ions, 
while  others,  like  alum,  decompose  in  aqueous  solution  into  their 
constituents  and  exist  only  in  the  solid  condition.  A  few  of  the 
more  important  double  salts  will  be  described  later. 

The  compounds  formed  by  the  union  of  salts  derived  from 
different  acids  are  called  mixed  salts;  those  which  contain  but  one 
acid  radical  and  decompose  largely  in  solution  into  the  simple 
salts  of  which  they  are  composed  are  called  double  salts;  and 
those  derived  from  a  single  acid  that  are  stable  in  solution  and 
yield  complex  ions  are  called  complex  salts. 

599.  The  lonization  of  Salts. — Although  the  acids  differ  widely 
in  the  extent  to  which  they  ionize,  the  salts  of  most  acids  are  highly 
dissociated  in  solution.     For  example,  sulphuric  acid  is  ionized  to 
the  extent  of  about  61  per  cent  in  a  one-tenth  normal  solution  at 
18°,  and  carbonic  acid  even  at  a  higher  dilution,  N/25,  is  but  0.2 
per  cent  ionized;   on  the  other  hand,  the  ionization  of  potassium 
sulphate  is  72  per  cent  and  that  of  potassium  carbonate  about  70 
per  cent.     The  ionization  of  acetic  acid  and  that  of  its  sodium  salt 
in  N/20  solution  are  2  per  cent  and  79  per  cent,  respectively. 

The  extent  to  which  a  salt  ionizes  is  determined  largely  by  the 
valence  of  the  ions  it  yields.  Salts  which  yield  two  univalent  ions 
are  the  most  highly  dissociated.  The  ionization  values  of  a  few 
salts  of  this  type  at  18°  in  one-tenth  normal  solution  are  as  follows: 
sodium  chloride  84  per  cent,  potassium  nitrate  83  per  cent,  silver 
nitrate  81  per  cent,  and  potassium  chlorate  83  per  cent.  If  one 
of  the  ions  is  univalent  and  the  other  bivalent  the  dissociation  is 
less.  Examples  of  salts  of  this  class  are  potassium  sulphate  72 
per  cent,  sodium  sulphate  70  per  cent,  disodium  phosphate  73 
per  cent,  barium  chloride  77  per  cent,  and  zinc  chloride  73  per  cent. 
If  both  ions  are  bivalent  the  ionization  is  still  less;  copper  sulphate 
39  per  cent,  zinc  sulphate  40  per  cent. 

600.  The  Hydrolysis  of  Salts.— The  fact  that  the  salts  of  weak 
acids  or  bases  are  hydrolyzed  by  water  has  been  frequently  men- 
tioned.    It  will  be  recalled  that  the  hydrolysis  of  salts  is  the  oppo- 
site of  neutralization.     It  is  advisable  at  this  point  to  review 
reactions  of  this  kind  in  the  light  of  what  has  been   recently 
learned. 

Neutralization  of  an  acid  by  a  base  is  brought  about  as  the 
result  of  the  union  of  hydrogen  and  hydroxyl  ions.  Water  itself 


512  INORGANIC  CHEMISTRY  FOR  COLLEGES 

is  ionized  to  the  extent  of  0.00001  per  cent,  and  an  equilibrium 
exists  between  the  molecules  and  the  ions  produced: 

0.00001% 
H2O     ;=±    H+  +  OH~ 

When  hydrogen  and  hydroxyl  ions  are  brought  together  by  adding 
a  solution  of  a  base  to  one  of  an  acid,  the  ions  react  to  form  water 
until  this  equilibrium  is  established.  The  reaction  between 
sodium  hydroxide  and  hydrochloric  acid  can  be  formulated  as 
follows : 

Na+  +  OH-  +  H+  +  OP  —  Na+  +  H20  +  Cr 

Since  sodium  and  chlorine  ions  do  not  take  part  in  the  reaction  if 
the  solutions  are  sufficiently  dilute,  the  equilibrium  is  between 
H+  and  OH~  on  one  side  and  H2O  on  the  other,  and  the  product 
of  the  concentration  of  these  ions  is  the  same  as  in  pure  water. 

When  the  salt  of  a  weak  acid  is  dissolved  in  water,  hydrolysis 
takes  place,  that  is,  a  part  of  the  salt  is  converted  into  acid  and 
base  by  the  water.  This  can  be  illustrated  by  a  consideration  of 
the  hydrolysis  of  sodium  hypochlorite.  Hypochlorous  acid  is 
dissociated  in  one-tenth  normal  solution  at  18°  to  the  extent  of 
0.02  per  cent.  There  is,  accordingly,  an  equilibrium  established 
between  the  acid  and  its  ions: 

0.02% 

HCIO  ^  H+  +  cicr 

If  C10~  ions  and  H+  ions  are  present  in  the  same  solution  this 
equilibrium  must  be  established.  This  occurs  when  sodium  hypo- 
chlorite is  dissolved  in  water,  for  the  former  furnishes  C1O~~  ions 
and  the  latter  H+  ions.  As  a  consequence,  under  these  conditions 
three  equilibria  are  set  up  which  may  be  represented  as  follows: 

H20  ^  H+  +  OH~ 

+ 
NaOCl  ^  C1CT  +  Na+ 

11 
HCIO 

Some  hydrogen  ions  from  the  water  unite  with  the  C10~  ions  to 
form  HCIO,  and,  as  a  consequence,  an  equivalent  quantity  of 
OH~  ions  are  left  in  the  solution,  which  shows,  therefore,  an  alka- 


THE  PROPERTIES  OF  OXIDES,  HYDROXIDES,  AND  SALTS     513 

line  reaction.  The  complete  reaction  may  be  represented  as 
follows : 

Na+  +  CIO"  +  H2O  ^  Na+  +  HC1O  +  OH~ 

Water  reacts  with  the  salt  to  form  the  base  and  acid,  but  as  the 
latter  is  only  slightly  dissociated  free  hydroxyl  ions  are  present 
in  the  solution  which  is,  accordingly,  alkaline.  To  sum  up  the 
subject  to  this  point  briefly, — neutralization  results  from  the 
union  of  H+  +  OH~  to  form  undissociated  water,  and  hydrolysis 
results  from  the  union  of  H+  and  an  acid  radical  to  form  molecules 
of  the  undissociated  acid. 

The  extent  to  which  a  salt  of  a  weak  acid  is  hydrolyzed  is  deter- 
mined by  the  strength  of  the  acid;  the  weaker  the  acid  the  greater 
the  hydrolysis,  because  more  hydrogen  ions  will  be  required  to 
convert  the  negative  ion  into  the  undissociated  acid,  and  more 
hydroxyl  ions  will  be  left  free  in  the  solution. 

A  similar  explanation  can  be  given  of  the  hydrolysis  of  salts 
derived  from  weak  bases:  in  this  case  the  undissociated  base  is 
formed  and  free  hydrogen  ions  are  left  in  solution.  In  the  case  of 
aluminium  sulphate,  for  example,  the  reaction  is  as  follows: 

A12(SO4)3  +  6H2O  ^  2A1(OH)3  +  3H2SO4 

If  the  salt  is  one  of  a  weak  acid  and  a  weak  base  the  extent  of 
the  hydrolysis  is  greater  because  there  are  two  active  agencies 
which  bring  it  about — the  formation  of  both  the  undissociated 
base  and  the  undissociated  acid. 

The  degree  of  hydrolysis  is  affected  not  only  by  the  strength  of 
the  acid  and  base  involved,  but  also  by  the  concentration  of  the 
solution — the  more  water  the  greater  the  hydrolysis — and  by  the 
temperature.  For  example,  copper  sulphate  is  hydrolyzed  suffi- 
ciently at  room-temperature  to  show  an  acid  reaction  with 
litmus  paper,  but  the  hydroxide  of  the  metal  is  not  formed  in 
large  enough  quantity  to  precipitate.  If  a  dilute  solution  of  the 
salt  is  boiled,  more  extensive  hydrolysis  takes  place  and  a  basic 
salt  formed  as  the  result  of  the  partial  hydrolysis  of  the  neutral 
salt  is  slowly  precipitated.  A  solution  of  magnesium  chloride  at 
100°  is  sufficiently  hydrolyzed  to  produce  a  concentration  of 
hydrochloric  acid  which  will  attack  iron;  it  is  for  this  reason  that 
water  to  be  used  in  boilers  should  be  freed  from  this  salt. 


514  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  relation  between  the  degree  of  hydrolysis  and  the  strength 
of  acids  and  bases  is  indicated  by  the  following  examples:  Borax 
is  hydrolyzed  to  the  extent  of  about  0.5  per  cent  in  one-tenth 
normal  solution;  the  ionization  of  boric  acid  at  this  concentration 
is  0.01  per  cent.  Aluminium  sulphate  in  0.001  N  solution  is 
hydrolyzed  to  the  extent  of  4.5  per  cent  at  25°.  Aluminium  car- 
bonate and  sulphide  are  both  completely  hydrolyzed  by  water. 

EXERCISES 

1.  Draw  up  a  general  statement  of  the  physical  and  chemical  properties 
of  the  hydroxides  of  metals  having  the  valence  1,  2,  and  3. 

2.  If  the  hydroxide  of  a  metal  has  amphoteric  properties  would  you  expect 
it  to  be  a  strong  base  or  a  strong  acid? 

3.  If  you  were  given  a  sample  of  a  chemical  element  how  could  you  deter- 
mine by    (a)  physical  and    (6)  chemical  means  whether  it  was  metallic  or 
non-metallic  in  character? 

4.  Give  examples  of  the  effect  of  change  in  valence  of    (a)  acid-forming 
elements  and    (6)  base-forming  elements  on  the  chemical  properties  of  their 
compounds  containing  oxygen  and  hydrogen. 

5.  Would  you  expect  that  the  chloride  of  an  acid-forming  element  would 
be  hydrolyzed  by  water?     Give  several  examples. 

6.  From  the  facts  given  in  this  chapter  state  how  you  could  distinguish 
the  following  from  each  other:    (a)  Fe(OH)2  and  Fe(OH)3,    (6)  Ca(OH)2, 
and   Zn(OH)2   (c)  A12(SO4)3  and  ZnSO4,    (d)  Zn(OH)2  and    Mn(OH)2,     (e) 
Cu(OH)2  and  Fe(OH),,    (/)  AgCl  and  PbCl,,    (g)  Na2CO3  and  NaCl. 

7.  Show  by  chemical  equations,  using  ionic  symbols,  what  happens  when 
ammonium  hydroxide  is  added  to  a  solution  of  calcium  phosphate  in  nitric 
acid. 

8.  (a)  How  do  you  think  the  two  oxides  of  tin,  SnO  and  SnO2,  would 
differ  in  chemical  properties?     (6)  Make  a  statement  as  to  the  behavior  of 
SnCl2  and  SnCl4  with  water. 

9.  Compare  the  properties  of  the  compounds  which  contain    (a)  anti- 
mony with  the  valence  3  and   (6)  with  the  valence  5. 

10.  When  Cr2Os  is  heated  in  the  air  it  is  stable.     If  it  is  mixed  with  solid 
sodium  hydroxide  and  heated  in  the  air  it  is  converted  into  sodium  chromate, 
Na2CrO4,  in  which  chromium  has  the  valence  6.     Can  you  explain  why  the 
reaction  takes  place? 


CHAPTER  XXXV 
SODIUM,  POTASSIUM,  RUBIDIUM,  AND  CAESIUM 

601.  Compounds  of  sodium  and  potassium  are  widely  dis- 
tributed, play  an  important  part  in  natural  processes,  and  are  used 
extensively  in  the  arts.  Lithium,  rubidium,  and  caesium,  the 
other  members  of  the  first  family  of  the  first  group  in  the  periodic 
classification  of  the  elements,  occur  in  but  small  quantities  in 
nature  and  the  uses  to  which  they  have  been  put  are  very  limited. 

The  elements  of  this  family  are  the  most  active  of  the  metals; 
they  oxidize  quickly  in  the  air,  and  are,  therefore,  kept  under  oil; 
they  decompose  water  rapidly  and  form  soluble  hydroxides,  which 
are  strong  bases  and  are  called  caustic  alkalies.  Their  salts  are 
very  stable  and  resist  high  temperatures.  The  hydroxide  of  caesium 
is  the  strongest  base,  a  fact  in  accord  with  the  conclusion  already 
stated  that,  in  most  cases,  increase  in  atomic  weight  in  a  chem- 
ical family  is  associated  with  increase  in  basic  properties. 

The  relationships  between  the  chemical  and  physical  properties 
of  potassium,  rubidium,  and  caesium  and  their  compounds  are 
very  close,  and  there  is  a  gradation  in  these  properties  similar 
to  that  emphasized  in  the  case  of  the  halogens.  Lithium  and 
sodium,  however,  do  not  fit  so  well  into  the  family  from  this 
point  of  view;  the  divergence  here  is  similar  to  that  noted  in 
the  case  of  fluorine  and  the  other  members  of  the  halogen  family, 
and  nitrogen  and  the  phosphorus  family  (449).  The  melting- 
points  of  the  elements  decrease  from  lithium,  186°,  to  caesium, 
26.5°,  but  the  specific  gravity  of  sodium,  0.97,  is  greater  than  that 
of  both  lithium,  0.53,  and  potassium.  0.86.  The  chlorides  of 
potassium,  rubidium,  and  caesium  are  isomorphous  with  one 
another,  but  not  with  sodium  chloride.  Only  the  more  impor- 
tant compounds  of  sodium  and  potassium  will  be  considered 
below. 

515 


516 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


SODIUM 

602.  Sodium   occurs    in   sodium    chloride,    complex  silicates, 
borax,   Chile  saltpeter,  cryolite,   and  in  other  minerals.      It  is 
present  in  sea-plants  as  salts  of  organic  acids,  which  are  con- 
verted into  sodium  carbonate  when  the  plants  are  burned.     Sodium 
chloride  is  an  important  constituent  of  the  blood. 

Sodium  was  first  isolated  in  1807  by  Davy,  who  obtained  it 
by  the  action  of  an  electric  current  on  moist  sodium  hydroxide. 

The  metal  is  manufactured  by  a 
process  invented  by  Castner.  Fused 
sodium  hydroxide  is  electrolyzed  in 
an  iron  vessel  in  which  is  suspended 
the  anode,  a,  (Fig.  37).  Over  the 
cathode,  c,  is  placed  a  drum,  d,  the 
lower  part  of  which  is  made  of  iron 
gauze,  g.  Oxygen  is  evolved  at  the 
anode  and  rising  through  the  fused 
mass  escapes  at  e.  Sodium  and 
hydrogen  liberated  at  the  cathode 
collecting;  the  gas  and  the  molten 
metal  are  led  off  through  pipes.  As 
the  electrolysis  proceeds  fresh  alkali 
is  added  through  the  hole  at  e. 
When  the  cell  is  started  the  hydrox- 
ide is  fused  by  gas  jets  at  i,  but  later  the  passage  of  the  current 
furnishes  the  heat  required  to  keep  the  electrolyte  molten. 

The  chief  physical  and  chemical  properties  of  sodium  have 
already  been  given  (page  443) .  It  is  so  soft  it  can  be  readily 
cut  with  a  knife.  It  dissolves  in  liquid  ammonia,  and  in  mercury 
(671).  It  unites  with  the  non-metals,  and  forms  a  crystalline 
compound  with  hydrogen  at  365°,  sodium  hydride,  NaH,  which 
reacts  with  water  to  form  sodium  hydroxide  and  hydrogen. 

Sodium  is  used  in  the  preparation  of  sodium  peroxide,  sodium 
cyanide,  NaCN,  and  many  organic  compounds.  It  is  also  used 
in  the  laboratory  for  drying  liquids,  such  as  ether,  which  do  not 
react  with  it. 

603.  Sodium  Chloride. — Common  salt  and  Chile  saltpeter  are 
the  sources  from  which  all  the  compounds  of  sodium  are  made. 


FIG.  37. 


SODIUM,  POTASSIUM,  RUBIDIUM,  AND  CAESIUM      517 

The  nitrate  is  used  in  making  nitric  acid  and  the  sodium  sulphate 
formed  as  a  by-product  is  used  as  such  or  in  making  glass  and  other 
substances.  Salt,  on  account  of  its  cheapness,  is,  however,  the 
chief  source  of  sodium  compounds.  It  is  obtained  as  rock  salt 
and  from  sea-water,  or  from  salt-brines  derived  from  lakes  or  wells. 

Rock  salt  is  found,  at  times,  in  a  very  pure  condition  and,  after 
being  mined,  is  ground  and  used  without  purification.  It  occurs  in 
the  United  States,  England,  Austria,  Germany,  and  Spain  in  large 
deposits.  Rock  salt  is  usually  contaminated  with  oxides  of  iron, 
clay,  sand,  and  other  substances,  and  is  purified  by  crystallization. 
It  is  mined  in  New  York,  California,  Kansas,  Utah,  and  Louisiana. 

The  chief  sources  of  salt  in  the  United  States  are  brines  found 
in  New  York  near  Syracuse  and  Warsaw,  in  Michigan  at  Saginaw 
Bay  and  Manistee,  and  near  Salina,  Kansas.  The  brine  occurs 
some  distance  below  the  surface  and  is  obtained  by  pumping  from 
wells,  which  are  bored  about  8  inches  in  diameter  and  lined  with 
iron  casings. 

The  brine  is  treated  with  milk  of  lime  or  sodium  carbonate  to 
precipitate  calcium  and  magnesium  compounds;  it  is  then  sepa- 
rated and  evaporated  to  obtain  the  salt.  In  the  commercial 
process  complete  purification  is  not  effected,  and  as  a  consequence, 
ordinary  salt  is  more  or  less  deliquescent.  Sodium  bicarbonate  is 
sometimes  added  to  table-salt  to  convert  the  magnesium  chloride 
present  into  the  insoluble  carbonate,  which  does  not  absorb  water 
from  the  air.  Very  pure  sodium  chloride  may  be  prepared  for 
chemical  purposes  by  conducting  hydrogen  chloride  into  a  satu- 
rated solution  of  salt  until  no  more  of  the  gas  dissolves.  Salt  is 
practically  insoluble  in  concentrated  hydrochloric  acid,  and  is 
precipitated,  while  the  impurities  remain  in  solution. 

Salt  crystallizes  in  cubes,  the  faces  of  which  are  usually  hollow. 
As  the  faces  grow,  layer  by  layer,  they  include  mechanically 
between  them  a  small  amount  of  the  liquid  from  which  they  are 
formed.  As  a  consequence,  when  the  crystals  are  heated  to  a  high 
temperature  the  liquid  is  vaporized  and  the  steam  produced  causes 
them  to  fly  apart  with  a  crackling  noise;  the  salt  decrepitates. 

Salt  melts  at  £20°  and  boils  at  1750°.  Its  solubility  is  given 
in  the  diagram  on  page  506.  Salt  is  a  necessary  constituent  of  the 
diet  of  animals,  as  it  plays  an  important  part  in  the  blood;  it  is 
used  as  the  source  of  chlorine  and  hydrochloric  acid,  in  the  man- 


518  INORGANIC  CHEMISTRY  FOR  COLLEGES 

ufacture  of  sodium  compounds,  in  the  preservation  of  meat  and 
fish,  in  glazing  pottery,  and  for  other  purposes. 

604.  Sodium  Hydroxide. — This  compound,  which  is  commonly 
called  caustic  soda  in  trade,  is  manufactured  by  the  electrolysis 
of  sodium  chloride  or  by  treating  a  solution  of  sodium  carbonate 
("  soda-ash  ")  with  milk  of  lime,  which  is  a  suspension  in  water  of 
finely  divided  calcium  hydroxide: 

Na2CO3  +  Ca(OH)2  =  2NaOH  +  CaCO3 

Steam  is  blown  into  the  mixture  in  order  to  agitate  the  solid  and 
thus  keep  it  in  contact  with  the  solution.  When  the  reaction  is 
complete,  the  calcium  carbonate  is  allowed  to  settle,  and  the  dilute 
solution  of  sodium  hydroxide  drawn  off  and  evaporated,  first  in 
vacuum  kettles  until  the  impurities  crystallize  out,  and  then  in  iron 
pots  over  an  open  fire.  When  all  the  water  has  been  driven  off, 
the  fused  sodium  hydroxide  is  poured  into  iron  drums. 

Large  quantities  of  caustic  soda  are  manufactured  by  the  elec- 
trolysis of  a  solution  of  sodium  chloride.  The  process  involves 
the  use  of  electrical  energy,  but  the  cost  of  this  is  offset,  in  part, 
by  the  fact  that  sodium  chloride  is  used  rather  than  sodium  car- 
bonate; the  latter  must  first  be  manufactured  from  salt  and  is, 
therefore,  the  more  expensive  substance;  and,  further,  the  chlorine 
formed  at  the  same  time  is  a  valuable  by-product. 

605.  When  an  electric  current  is  passed  through  a  solution  of 
sodium  chloride,  sodium  hydroxide  and  hydrogen  are  produced 
at  the  cathode  and  chlorine  at  the  anode.     If  the  two  portions  of 
the  solution  are  allowed  to  mix,  the  halogen  and  alkali  react  and 
sodium  hypochlorite  is  formed.     The  large  number  of  cells  which 
have  been  invented  to  prepare  sodium  hydroxide  and  chlorine 
electrolytically  differ  from  one  another  chiefly  in  the  way  in  which 
the  formation  of  the  hypochlorite  is  avoided.     In  one  type  of 
cell  a  diaphragm  is  used  to  separate  the  liquid  around  the  cathode 
from  that  around  the  anode.     The  diaphragm  is  constructed  of 
porous  material,  which  allows  the  liquid  to  penetrate  it,  and  thus 
does  not  prevent  the  flow  of  the  electric  current;  but  it  does  pre- 
vent mechanical  mixing  of  the  solutions  in  the  two  compartments. 
One  form  of  diaphragm  used  is  made  of  a  mixture  of  asbestos  and 
iron  oxide  supported  on  an  iron  grating,  which  serves  as  the  cathode. 
The  Townsend  cell,  which  is  an  example  of  this  type,  is  represented 


SODIUM,  POTASSIUM,  RUBIDIUM,  AND  CAESIUM       519 


sHydrogen 
'*    Outlet 


MiL%5: 


LJ 


-JBfmeJ\ 


Lye  Outlet 

i 

'Connection 
to  Cathode 


diagrammatically  in  Fig.  38.  A  is  the  anode,  which  is  made  of 
graphite.  A  solution  of  a  chloride  is  allowed  to  flow  slowly  into 
the  cell,  and  is  kept  at  the  level  indicated.  The  cathode  is  of 
perforated  iron  and  supports  the  diaphragm.  Kerosene  is  placed 
in  the  outer  compartment.  /+ 

When  the  electrolysis  is  taking  nn    ..-Anode 

place  the  solution  slowly  per- 
colates through  the  diaphragm 
due  to  the  hydrostatic  pressure 
of  the  water;  it  falls  in  drops 
through  the  kerosene  to  the  V 

bottom  of  the  outer  compart- 
ment from  which  it  passes  out 
through  a  pipe.  The  solution 
obtained  contains  about  150 
grams  of  sodium  hydroxide  and 
200  grams  of  sodium  chloride  FIG.  38. 

per  liter.    It  is  evaporated  until 

the  salt  crystallizes  out,  separated  from  the  solid,  and  finally 
heated  in  iron  pots  until  all  the  water  has  been  driven  off.  The 
molten  sodium  hydroxide  is  then  poured  into  iron  containers. 

The  Allen-Moore  and  Nelson  cells  resemble  in  construction 
the  Townsend  cell,  but  kerosene  is  not  used  in  the  outer  com- 
partments. 

606.  The  Castner-Kellner  cell  (Fig.  39)  is  an  example  of  a 
different  type,  in  which  no  diaphragm  is  used.  The  anodes  are 


FIG.  39. 

of  graphite,  and  are  placed  in  compartments  separated  by  par- 
titions from  the  compartment  containing  the  cathode,  which  is 
made  of  iron.  A  layer  of  mercury  is  placed  on  the  bottom  of 


520  INORGANIC  CHEMISTRY  FOR  COLLEGES 

the  cell  as  indicated.  The  chlorine  liberated  escapes  through  the 
outlet  indicated  in  the  diagram.  Metallic  sodium  is  set  free  at 
the  surface  of  the  mercury  in  the  anode  compartments  and 
dissolves  in  the  mercury.  During  the  electrolysis  the  cell  is 
rocked  slowly  by  an  eccentric,  and  the  mercury  flows  from  the 
anode  to  the  cathode  compartment  and  back  again.  The  depres- 
sions at  the  bottom  of  the  tank  are  of  such  a  size  that  during  the 
rocking  none  of  the  solution  from  the  anode  compartment  gets 
into  the  cathode  compartment.  When  the  mercury  containing 
the  sodium  in  solution  reaches  the  cathode  compartment,  the 
sodium  reacts  with  water  to  form  sodium  hydroxide,  and  hydrogen 
is  evolved  at  the  cathode.  A  strong  solution  of  sodium  hydroxide 
free  from  salt  is  obtained. 

607.  Sodium  hydroxide  is  a  white  solid  which  absorbs  water 
from  the  air,  and  serves,  therefore,  as  a  valuable  drying  agent. 
It  melts  at  red-heat.     It  is  used  in  large  quantities  in  the  man- 
ufacture of  soap,  paper-pulp,  and  phenol,  and  other  important 
organic  compounds  used  in  the  preparation  of  dyes. 

608.  Sodium  Carbonate. — The  mineral  called  trona,  which  has 
the  composition  NaHCO3,Na2CO3,2H2O,  and  the  ashes  from  sea- 
plants  were  used  as  the  sources  of  sodium  carbonate  until  the 
development  of  the  chemical  industries  made  it  imperative  to 
invent  a  process  for  the  preparation  of  this  very  important  com- 
pound from  salt.     As  the  result  of  war  France  was  cut  off  from  a 
supply  of  the  carbonate  and  a  prize  of  100,000  francs  was  offered 
by  the  French  Academy  for  the  best  solution  of  the  problem. 
Le  Blanc  invented  a  process,  which  has  since  been  known  by  his 
name,  won  the  prize,  and  was  granted  a  patent  in  1791.     During 
the  French  Revolution  his  factory  was  seized  and  declared  to  be 
public  property.     He  was  paid  no  indemnity  and,  finally,  in 
discouragement,  committed  suicide. 

The  process,  with  but  slight  improvement,  is  used  extensively 
to-day,  although  in  the  United  States  it  has  been  superseded  by 
another  process  invented  by  Solvay,  a  Belgian  chemical  manu- 
facturer. In  both  cases  sodium  chloride  is  the  source  of  the 
sodium,  and  calcium  carbonate  is  the  source  of  the  carbonate 
radical,  but  the  reactions  used  to  effect  the  formation  of  sodium 
carbonate  are  different.  The  value  of  the  by-products  produced 
largely  determines  the  net  cost  of  the  carbonate  and,  consequently, 


SODIUM,  POTASSIUM,  RUBIDIUM,  AND  CAESIUM       521 

the  process  to  be  used  under  the  conditions  which  exist  in  the 
chemical  market.  It  will  be  of  interest,  therefore,  to  consider  in 
some  detail  these  two  processes  to  see  how  the  costs  of  materials 
and  energy  and  the  value  of  by-products  are  the  most  important 
factors  in  industrial  chemistry  from  the  economic  point  of  view. 

609.  The  Le  Blanc  Process. — It  will  be  recalled  that  when  a 
solution  of  sodium  carbonate  is  added  to  one  of  calcium  chloride  a 
double  decomposition  takes  place  owing  to  the  insolubility  of 
calcium  carbonate; 

Na2CO3  +  CaCl2  =  CaC03  +  2NaCl 

Since  sodium  chloride  does  not  interact  with  calcium  carbonate, 
the  reaction  cannot  be  made  to  proceed  in  the  reverse  direction. 
Le  Blanc's  solution  of  the  problem  was  to  convert  the  chloride 
into  another  salt,  which,  when  fused  with  calcium  carbonate, 
entered  into  double  decomposition  with  it.  Sodium  sulphide 
acts  in  this  way: 

Na2S  +  CaCOs  =  CaS  +  Na2CO3 

The  sulphide  was  prepared  by  converting  the  chloride  into  the 
sulphate, 

2NaCl  +  H2SO4  =  Na2SO4  -f  2HC1, 

and  fusing  the  latter  with  carbon: 

Na2SO4  +  2C  =  Na2S  +  2CO2 

The  products  used  are,  accordingly,  salt,  coal,  calcium  carbonate, 
and  sulphuric  acid;  the  by-products  are  hydrochloric  acid,  cal- 
cium sulphide,  and  carbon  dioxide.  The  economic  success  of  such 
a  process  depends  upon  the  utilization  of  these  by-products;  sul- 
phuric acid  must  be  manufactured  and  is,  thus,  an  important  item 
in  the  expense.  At  first  the  hydrochloric  acid  formed  was  a  nui- 
sance, for  it  was  discharged  into  the  air,  but  as  the  chemical  indus- 
tries developed  it  found  important  uses  and  became,  therefore,  of 
value.  The  acid  not  required  for  these  purposes  is  converted 
into  chlorine,  which  is  largely  used  to  make  bleaching  powder. 
In  this  way  this  by-product  becomes  of  value,  and  serves  to 
reduce  the  cost  of  the  sodium  carbonate. 


522  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  sulphur  in  the  calcium  sulphide  is  recovered  by  what  is 
known  as  the  Chance-Glaus  process.  The  sulphide  is  suspended 
in  water  and  treated  with  carbon  dioxide,  which  sets  free  hydrogen 
sulphide : 

CaS  +  H20  +  C02  =  CaC03  +  H2S 

The  latter  is  then  burned  in  a  furnace  with  just  enough  air  to  con- 
vert it  into  sulphur: 

2H2S  +  O2  =  2H20  +  2S 

The  sulphur  is  sold  or  made  into  sulphuric  acid  to  be  used  in  the 
process.  About  85  per  cent  of  the  sulphur  in  the  sulphuric  acid 
used  in  converting  the  salt  into  sodium  sulphide  can  be  recovered 
by  this  process. 

610.  The  Solvay  Process. — The  reactions  upon  which  this  proc- 
ess is  based  were  discovered  in  1838.  Attempts  were  made  early 
to  utilize  the  reactions  for  industrial  purposes,  but  they  were 
unsuccessful,  largely  because  the  process  involved  the  handling  of 
gases  on  the  large  scale  and  the  careful  regulation  of  temperature, 
the  technique  of  which  had  not  been  fully  worked  out.  Solvay 
proved  himself  to  be  a  capable  engineer  and  overcame  the  difficul- 
ties. In  1863  he  designed  the  necessary  apparatus  for  a  plant 
and  in  1873  the  process  had  established  itself  as  a  commercial 
success.  It  is  the  one  used  in  the  United  States  for  the  manu- 
facture of  sodium  carbonate. 

In  this  process  sodium  chloride  is  first  converted  into  sodium 
bicarbonate,  which  is  difficultly  soluble  in  water,  by  treating  it 
with  ammonium  bicarbonate,  which  is  soluble: 

NaCl  +  NH4HCO3  =  NaHCO3  +  NH4C1 
When  this  salt  is  heated  it  decomposes: 

2NaHCO3  =  Na2CO3  +  H2O  +  CO2 

The  ammonium  bicarbonate  required  is  made  by  the  action  of  car- 
bon dioxide  on  a  solution  of  ammonia  in  water: 

NH4OH  +  H2C03  =  NH4HC03  +  H20 

The  carbon  dioxide  is  obtained  by  heating  limestone,  CaCOs  = 
CaO  +  CO2 ;  and  the  lime  formed  is  used  to  recover  the  ammonia 


SODIUM,  POTASSIUM,  RUBIDIUM,  AND  CAESIUM      523 

from  the  ammonium  chloride  produced  according  to  the  first 
reaction  given  above: 

2NH4C1  +  Ca(OH)2  =  2NH3  +  2H20  +  CaCl2 

The  products  used  in  the  process  are,  accordingly,  sodium 
chloride  and  calcium  carbonate,  and  the  by-product  is  calcium 
chloride;  the  ammonia  is  recovered  and  used  over  and  over  again. 
It  is  seen,  thus,  that  through  the  ingenious  use  of  the  fact  that 
sodium  bicarbonate  is  difficultly  soluble  in  water,  it  is  possible  to 
realize  the  transformation  indicated  by  the  equation 

CaCO3  +  2NaCl  =  Na2C03  +  CaCl2 

which  is  the  reverse  of  that  which  expresses  the  facts  when  sodium 
carbonate  and  calcium  chloride  interact  in  solution. 

The  by-product  of  the  Solvay  process,  calcium  chloride,  is  of 
little  value;  it  is  cheap  enough  to  warrant  its  use  on  roads  to  lay 
dust.  This  method  of  preparing  sodium  carbonate  does  not 
yield  as  by-products  hydrochloric  acid  and  chlorine  as  the  Le  Blanc 
process  does.  Chlorine  is  made  in  the  United  States  by  the 
electrolysis  of  salt,  because  water-power  is  available  and  electrical 
energy  correspondingly  cheap.  It  is  this  factor  which  determines 
the  use  of  the  Solvay  process  in  America  and  that  of  Le  Blanc  in 
England. 

611.  Properties  of  Sodium  Carbonate  and  Sodium  Bicarbonate. 
— Sodium  carbonate  crystallizes  from  water  below  35.2°  as  a 
decahydrate,  Na2CO3,10H2O,  which  is  sometimes  called  washing 
soda  or  sal  soda.  It  effloresces  rapidly  in  the  air  and  is  converted 
into  a  white  powder  which  has  the  composition  of  the  crystals 
formed  from  solutions  above  35.2°,  namely,  Na2CO3,H2O.  The 
anhydrous  salt  is  called  in  trade  soda-ash. 

Sodium  carbonate  is  hydrolyzed  in  aqueous  solution,  3.2  per 
cent  being  converted  into  the  hydroxide  in  N/25  solution  at  25°: 

Na2CO3  +  H2O  ;=±  NaHCO3  +  NaOH 

Many  of  the  uses  of  sodium  carbonate  are  based  on  this  fact. 
It  is  used  in  washing  and  for  scouring  purposes  because  the  small 
amount  of  alkali  present  has  a  marked  effect  on  oil  or  grease;  it 
converts  them  into  the  colloidal  condition  as  the  result  of  the 


524  INORGANIC  CHEMISTRY  FOR  COLLEGES 

formation  of  minute  globules  which  become  suspended  in  the 
solution  and  are,  as  a  result,  easily  removed. 

Sodium  carbonate  is  used  in  large  quantities  in  softening  water 
(629),  in  the  manufacture  of  glass  and  soap,  and  in  smaller  amounts 
in  many  other  industries. 

Sodium  bicarbonate,  NaHCOs,  is  an  anhydrous  salt;  when  a 
solution  of  the  salt  in  water  is  boiled,  it  is  hydrolyzed  and  carbon 
dioxide  slowly  escapes;  sodium  carbonate  is  formed  to  some 
extent  and  owing  to  the  hydrolysis  of  the  latter  the  solution 
becomes  alkaline.  Sodium  bicarbonate  has  long  been  used  under 
the  name  baking  soda  or  saleratus.  In  making  leavened  bread 
it  is  necessary  to  add  to  the  dough  something  that  produces  bub- 
bles of  gas  which  expand  when  the  dough  is  baked,  and  thus  give 
to  the  bread  a  porous  structure.  Carbon  dioxide  is  the  gas  com- 
monly used  for  this  purpose;  it  is  generated  either  through  the 
use  of  yeast,  which  is  a  plant  that  evolves  carbon  dioxide  as  it 
grows,  or  as  the  result  of  the  interaction  of  sodium  bicarbonate 
and  an  acid  put  into  the  dough.  Molasses  and  vinegar  contain 
acids  and  can  be  used  with  the  salt,  or  cream  of  tartar,  acid 
potassium  tartrate,  KH(C4H4O6),  may  be  employed. 

612.  Baking  powders  are  mixtures  of  sodium  bicarbonate  and 
a  substance  which  reacts  with  the  latter  to  form  carbon  dioxide. 
The  tartrate  powders  contain  tartaric  acid,  H2(C4H4O6),  or  cream 
of  tartar,  the  phosphate  powders  contain  primary  calcium  or 
sodium  phosphate,  and  the  alum  powders  contain  aluminium 
sulphate  or  a  double  salt  of  the  latter  and  sodium  sulphate.  The 
reaction  which  takes  place  in  the  case  of  cream  of  tartar  powders 
is  as  follows: 

NaHC03  +  KH(C4H4O6)  -  KNa(C4H4O6)  +  H2O  +  CO2 

When  aluminium  sulphate  is  used,  the  salt  first  hydrolyzes  to  form 
aluminium  hydroxide  and  sulphuric  acid: 

A12(SO4)3  +  6H2O  =  2A1(OH)3  +  3H2SO4 
and  the  latter  decomposes  the  bicarbonate: 

2NaHCO3  +  H2SO4  =  Na2SO4  +  2H2O  +  2CO2 

The  products  left  in  the  bread  are,  accordingly,  either  potassium 
sodium  tartrate,  which  is  commonly  called  rochelle  salt,  or  alumin- 


SODIUM,  POTASSIUM,  RUBIDIUM,  AND  CAESIUM       525 

ium  hydroxide,  or,  in  the  case  of  the  phosphate  powders,  second- 
ary salts  of  the  acid. 

Starch  is  added  to  baking  powder  to  cover  the  grains  of  the 
chemicals  used  so  that  they  cannot  interact  until  the  powder  is 
moistened  in  the  dough.  It  is  evident  that  baking  powders  should 
be  protected  from  the  air  which  contains  moisture.  The  starch 
also  serves  as  a  diluent,  and  the  amount  added  to  each  kind  of 
powder  is  that  required  to  give  a  mixture  which  evolves  about  12 
per  cent  of  its  weight  of  carbon  dioxide.  As  a  result,  the  same 
amount  of  different  powders  have  about  the  same  leavening  effect. 
Sodium  bicarbonate  is  used  in  Seidlitz  powders,  bromo-seltzer, 
and  other  effervescing  medicinal  drinks. 

613.  Other  Salts  of  Sodium. — A  number  of  sodium  salts  have 
already  been  described  more  or  less  fully.     Among  these  are  the 
following:    Borax,   sodium  nitrite,   sodium  nitrate,   sodium  sul- 
phate,   sodium   thiosulphate,    sodium   sulphite,    sodium   silicate, 
sodium  peroxide,  and  the  sodium  phosphates. 

614.  The  Test  for  Sodium  Salts.— Two  difficultly  soluble  salts 
are  used  at  times  in  qualitative  tests  for  the  element;   these  are 
sodium  fluosilicate,  Na^SiFe,  which  is  formed  by  adding  hydro- 
fluosilicic  acid  to  a  strong  solution  of  a  sodium  salt,  and  sodium- 
hydrogen  pyroantimonate,  Na2H2Sb2C>7,  which  is  formed  when  the 
potassium  salt  of  the  acid  is  added  to  the  solution.     The  test 
commonly  applied  is  the  examination  of  the  spectrum  of  the  light 
produced  when  some  of  the  solution  suspended  on  a  platinum  wire 
is  put  into  a  Bunsen  flame.     Sodium  compounds  give  a  bright 
yellow  flame.     If  the  characteristic  lines  persist  for  some  time  and 
are  not,  therefore,  produced  as  the  result  of  traces  of  sodium  com- 
pounds as  accidental  impurities,  the  presence  of  the  element  is 
shown.     The  compounds  of  several  elements  yield  colored  flames 
when  they  are  placed  in  a  Bunsen  flame.     In  order  to  determine 
what    lements  are  present,  the  light  is  examined  with  a  spectro- 
scope. 

In  the  quantitative  determination  of  sodium,  all  other  elements 
are  separated  and  the  metal  is  finally  converted  into  sodium 
sulphate,  which  is  heated  to  a  high  temperature  and  weighed. 

615.  Spectroscope. — When    a    beam    of   white    light    passes 
through  a  triangular  prism  made  of  glass,  it  is  broken  up  into  its 
constituents,  and  a  spectrum  is  obtained  which  shows  the  colors  of 


526  INORGANIC  CHEMISTRY  FOR  COLLEGES 

the  rainbow.  This  results  from  the  fact  that  white  light  is  made  up 
of  vibrations  of  different  wave-lengths,  each  of  which  produces  a 
characteristic  color.  If  the  light  produced  when  a  sodium  salt 
is  heated  in  a  Bunsen  flame  is  passed  through  a  slit  to  produce 
a  beam  and  then  through  the  prism,  a  bright  yellow  line  is 
obtained  and  the  rest  of  the  spectrum  is  lacking.  If  a  potassium 
salt  is  used  a  number  of  blue  lines  are  visible  along  with  a  few 
red  and  green  lines.  Each  element  in  the  free  condition  or  in 
combination,  when  it  is  heated  to  a  sufficiently  high  temperature, 
gives  in  a  similar  way  a  characteristic  spectrum  made  up  of 
colored  lines.  The  spectra  of  the  elements  have  been  determined, 
and  the  presence  of  any  one  of  the  latter  in  a  mixture  can  be  dis- 
covered by  an  examination  of  the  spectrum  produced  when  the 
material  containing  it  is  heated  to  such  a  temperature  that  the 
characteristic  vibrations  are  produced.  The  optical  instrument 
by  means  of  which  the  light  is  examined  is  called  a  spectroscope. 
It  consists  essentially  of  a  tube  furnished  at  one  end  with  a  slit 
through  which  the  light  passes,  and  at  the  other  end  of  which  is  a 
lens  to  convert  the  light  into  a  parallel  beam.  The  latter  then 
passes  through  a  triangular  prism  and  the  issuing  beam,  which  is 
spread  out  into  its  constituents,  is  examined  by  means  of  a  small 
telescope.  In  the  direct  vision  spectroscope  the  lenses  and  prisms 
are  so  arranged  that  but  one  tube  is  necessary. 

At  the  temperature  produced  in  a  Bunsen  flame  only  the  spectra 
of  the  metals  of  the  alkalies  or  alkaline  earths  can  be  observed. 
The  chlorides  of  the  metals  are  commonly  used  on  account  of  their 
volatility. 

POTASSIUM 

616.  The  occurrence  of  potassium  in  feldspar  and  other  complex 
silicates  has  already  been  noted  (556).  The  chief  sources  of  the 
element  and  its  compounds  are  the  chloride  and  sulphate  present 
in  the  salt  deposits  in  Alsace  and  Stassfurt,  Germany,  where  they 
occur  as  sylvite,  KC1,  carnalite,  KCl,MgCl2,6H2O,  and  kainite, 
K2SO4,MgSO4,MgCl2,6H2O.  Smaller  quantities  are  obtained 
from  wood  ashes,  the  residue  from  beet  sugar,  and  wool  washings. 
Largely  as  the  result  of  the  recent  war  the  enormous  beds  of  sea- 
weed along  the  coast  of  California  have  been  used  as  a  source  of 
potassium  and  of  iodine.  The  ashes  obtained  by  burning  the  sea- 


SODIUM,  POTASSIUM,  RUBIDIUM,  AND  CAESIUM      527 

weed  contain  about  9  per  cent  potassium  chloride  and  0.1  per 
cent  iodine  as  potassium  iodide.  Large  deposits  of  alunite, 
K2SO4,A12(S04)3,4A1(OH)3,  exist  in  Utah,  Nevada,  and  Colorado, 
and  are  now  utilized  as  a  source  of  potassium  sulphate.  When 
the  mineral  is  roasted  the  sulphate  and  hydroxide  of  aluminium  are 
decomposed  into  oxide,  and  the  product  on  treatment  with  water 
yields  a  solution  of  potassium  sulphate.  Potassium  chloride  has 
recently  been  recovered  successfully  from  the  fumes  from  cement 
plants  and  blast  furnaces.  Although  much  attention  has  been 
paid  as  a  result  of  the  war  to  the  production  of  potassium  com- 
pounds from  the  natural  resources  of  the  United  States,  it  is  highly 
probable  that  Europe  will  for  some  time  supply  the  world  with 
potash  salts,  which  are  essential  plant-foods,  and,  therefore, 
important  constituents  of  fertilizers. 

Potassium  was  first  isolated  by  Davy  in  1807;  it  is  prepared 
by  a  method  similar  to  that  used  in  the  case  of  sodium.  Its  physical 
and  chemical  properties  resemble  those  of  sodium  (602). 

617.  Potassium  hydroxide  is  made  by  the  electrolysis  of  potas- 
sium chloride,  and  resembles  closely  sodium  hydroxide.    When  it  is 
required  in  a  very  pure  form  it  is  dissolved  in  alcohol  in  which 
the  chloride  and  carbonate,  the  common  impurities  present,  are 
insoluble.     After  the  clear  liquid  has  been  separated  from  the  in- 
soluble material,  it  is  evaporated  to  dry  ness.     Potassium  hydroxide 
is  very  deliquescent  and  very  soluble  in  water,  1  liter  of  which  will 
dissolve  about  1500  grams  of  the  solid.     It  is  called  caustic  potash 
in  trade.     It,  like  sodium  hydroxide,  decomposes  proteins  and  dis- 
solves the  flesh  if  left  in  contact  with  it. 

618.  Salts  of  Potassium. — Potassium  carbonate  cannot  be  made 
from  potassium  chloride  by  the  Solvay  process  because  potassium 
bicarbonate  is  readily  soluble  in  water.     It  is  for  this  reason  that  a 
solution  of  potassium  hydroxide  is  preferred  to  one  of  sodium 
hydroxide  as  an  absorbent  for  carbon  dioxide.     When  the  sodium 
hydroxide  is  used,  the  bicarbonate  separates  after  a  time  and 
clogs  up  the  tubes,  etc.,  used  in  the  absorbing  apparatus. 

The  carbonate  is  made  by  the  Le  Blanc  process  or  in  other  ways. 
Its  trade  name  is  potash  or  pearlash.  It  is  prepared  in  appre- 
ciable quantities  from  suint,  the  fatty  material  present  in  wool, 
which  contains  along  with  grease  the  potassium  salt  of  an  organic 
acid.  The  grease  is  extracted  and  purified  and  sold  under  the 


528  INORGANIC  CHEMISTRY  FOR  COLLEGES 

name  "  lanolin,"  and  the  residue  is  ignited  and  converted  into 
potassium  carbonate.  The  salt  is  usually  sold  in  the  anhydrous 
form,  and  differs  from  sodium  carbonate  in  being  very  deliquescent. 
The  crystals  obtained  from  water  have  the  composition  2K2COs,- 
3H2O.  It  is  used  in  the  manufacture  of  glass,  soft  soap,  and  certain 
potassium  salts. 

619.  Potassium  nitrate  was  formerly  prepared  from  the  products 
resulting  from  the  decomposition  of  animal  refuse  in  natural  or 
artificial  nitrate  beds,  but  during  the  Crimean  war  (1852-55)  the 
supply  of  this  important  ingredient  of  gunpowder  became  deficient 
and  the  method  now  used  was  invented.     This  involves  the  treat- 
ment of  potassium  chloride  with  sodium  nitrate  in  aqueous  solu- 
tions : 

KC1  +  NaNO3  =  KNO3  +  NaCl 

The  reaction  has  already  been  explained  (592). 

Potassium  nitrate  (saltpeter  or  niter)  melts  at  339°  and  at 
higher  temperatures  is  reduced  to  potassium  nitrite. 

Potassium  nitrate  is  used  in  the  manufacture  of  gunpowder, 
matches,  and  fireworks,  the  use  depending  on  the  fact  that  the  salt 
gives  off  oxygen  freely  at  high  temperatures  and  is,  therefore,  an 
active  oxidizing  agent.  It  is  also  used  in  curing  meats;  it  prevents 
putrefaction  and  produces  the  deep  red  color  familiar  in  the  case 
of  salted  hams  and  corned  beef. 

620.  Gunpowder  is  the  oldest  explosive  and  is  still  a  very 
important  one.     According  to  tradition  it  was  invented  by  a  monk 
at  Freiburg,  Germany,  in  the  fourteenth  century;  it  was  first  used 
at  the  battle  of  Crecy  in  1346.     Gunpowder  is  a  mixture  of  potas- 
sium nitrate,  sulphur,  and  charcoal.     Its  explosive  properties  are 
due  to  the  fact  that  when  it  is  ignited  there  is  produced  a  large 
volume  of  gas,  which  consists  principally  of  nitrogen  and  carbon 
dioxide.     A  mixture  of  potassium  nitrate  and   charcoal  burns 
vigorously  when  ignited,  but  a  large  part  of  the  carbon  unites  with 
the  potassium  to  form  potassium  carbonate.     In  order  to  prevent 
this  and  thus  increase  the  proportion  of  gas  formed,  sulphur  is 
added  to  the  mixture  in  sufficient  quantity  to  unite  with  the  metal 
to  form  potassium  sulphide.     When  gunpowder  burns  in  the  open 
air  a  reaction  takes  place  which  is  indicated  by  the  following 
equation : 

2KNO3  +  3C  +  S  =  K2S  +  3CO2  +  N2 


SODIUM,  POTASSIUM,  RUBIDIUM,  AND  CAESIUM      529 

When  it  is  exploded  in  a  confined  space  the  reaction  is  much  more 
complex,  other  substances  in  addition  to  those  just  mentioned 
being  formed,  among  which  are  potassium  carbonate,  potassium 
sulphate,  and  carbon  monoxide.  The  products  consist  of  about 
43  per  cent  gases  and  57  per  cent  solids  by  weight. 

621.  Other  Salts  of  Potassium. — A  large  number  of  potassium 
salts  are  known,  and  many  are  used  in  the  laboratory  in  preference 
to  the  cheaper  sodium  salts  when  the  acid  radical  of  the  salt  is 
desired.     For  example,  the  iodide  commonly  used  is  potassium 
and  not  sodium  iodide.     The  more  expensive  salts  are  employed 
because  in  most  cases  they  are  less  soluble  in  water,  and,  therefore, 
more  easily  purified  by  crystallization,  and  because  they  are  more 
stable  in  the  air.     Certain  sodium  salts  are  hygroscopic  whereas 
the  corresponding  potassium  salts  are  not;    in  the  case  of  the 
carbonates,  however,  the  potassium  salt  is  hygroscopic  and,  as  a 
consequence,  sodium  carbonate  is  the  ordinary  laboratory  reagent. 
Potassium  cyanide,  KCN,  is  made  by  heating  potassium  ferro- 
cyanide   (751).     It  yields  hydrocyanic  acid  (prussic  acid),  b.p. 
20.5°,  when  treated  with  an  acid.     The  salt  and  the  acid  are 
active  poisons. 

622.  Tests  for  Potassium  Salts. — The  presence  of  potassium 
salts  is  tested  for  qualitatively  by  observing  the  flame  produced 
by  introducing  into  a  Bunsen  flame  some  of  the  salt  on  a  platinum 
wire.     If  a  spectroscope  is  not  available,  the  flame  is  examined 
by  looking  at  it  through  a  piece  of  blue  cobalt-glass,  which  allows 
the  violet  light  produced  by  potassium  to  pass  through  it  but 
cuts  off  the  light  from  sodium  and  other  metals. 

Potassium  is  separated  from  sodium  quantitatively  by  adding 
to  the  solution  of  its  salts  chloroplatinic  acid,  H^PtCle,  which 
precipitates  potassium  chloroplatinate,  K^PtCle.  Since  the  salt 
is  somewhat  soluble  in  cold  water,  it  is  precipitated  and  washed 
in  a  mixture  of  water  and  alcohol  in  which  it  is  less  soluble. 
Potassium  perchlorate  is  very  slightly  soluble  in  water  and  is 
also  used  in  the  quantitative  determination  of  potassium. 

EXERCISES 

1.  Starting  with  sodium  chloride  as  the  source  of  sodium,  indicate  by 
chemical  equations  by  what  reactions  the  following  compounds  can  be  pre- 
pared: (a)  NaNO,,"(6)  Na2B4O7,  (c)  Na2HPO4,  (d)  Na2S2O7,  (e)  Na2O2, 
(/)  Na2S203. 


530  INORGANIC  CHEMISTRY  FOR  COLLEGES 

2.  (a)  Why  can  NaCl  be  readily  removed  from  many  other  salts  by 
crystallization?     How  would  you  proceed  to  get    (6)  pure  NaCl,  and   (c) 
KNO3,  from  a  mixture  of  the  two  salts? 

3.  State   the   economic   advantages   of   combining   a   plant   for   making 
ammonia  by  the  HaBer  process  with  one  making  sodium  carbonate  by  the 
Solvay  process  provided  NH^Cl  is  used  instead  of  (NH4)2SO4  in  fertilizers. 
Consider  raw  materials  and  by-products. 

4.  Why  cannot  Na2CO3  replace  NaHCO3  in  baking  powders? 

5.  Calculate    (a)  the  weight  of  cream  of  tartar  required  to  react  with 
100  grams  NaHCO3  when  the  two  substances  are  mixed  to  prepare  a  baking 
powder.     Calculate    (6)  the  weight  of  CO2  given  off  by  the  mixture    (c)  the 
percentage  by  weight  of  CO2  given  off  by  the  mixture,  and    (d)  the  weight 
of  starch  that  must  be  added  so  that  the  resulting  mixture  of  the  three  ingredi- 
ents will  yield  12  per  cent  of  its  weight  of  CC>2. 

6.  Starting  with  alunite,  indicate  by  equations  by  what  chemical  reac- 
tions the   following  could  be  prepared:     (a)  K2SO4,     (6)  A1C13,     (c)  KC1, 
(d)  alum  (K2S04,  A12(SO4)3,  24H20). 


CHAPTER  XXXVI 
CALCIUM,  STRONTIUM,  BARIUM,  AND  RADIUM 

623.  Calcium,  strontium,  and  barium  form  a  typical  family  of 
elements  which  are  known  as  the  alkaline  earth  metals;  their 
atomic  weights  are,  respectively,  40.07,  87.63,  and  137.37.  They 
always  have  the  valence  2,  and  the  general  chemical  behavior  of 
their  compounds  is  that  shown  by  the  compounds  of  other  metals 
when  they  exhibit  this  valence.  We  shall  see  as  the  chemistry  of 
the  metallic  elements  is  developed,  that  the  valence  shown  by  a 
metal  is  a  most  important  factor  in  determining  the  behavior  of  its 
salts. 

Barium,  the  member  with  the  highest  atomic  weight  in  the 
family,  is  the  most  active  in  base-forming  properties.  Like  potas- 
sium and  sodium  it  forms  a  peroxide,  BaO2,  when  heated  in  the  air; 
its  hydroxide  is  decomposed  only  at  high  temperatures  into  oxide 
and  water;  its  nitrate  is  reduced  by  heat  to  a  nitrite,  which,  unlike 
the  nitrites  of  potassium  and  sodium,  can  be  converted  at  a  high 
temperature  into  the  oxide  and  oxides  of  nitrogen.  The  sulphate 
and  carbonate  of  barium  require  high  temperatures  to  convert 
them  into  oxides.  In  the  stability  of  its  compounds  toward  heat, 
barium  stands  below  potassium  and  sodium,  but  resembles  them 
closely.  The  compounds  of  strontium  and  calcium  behave  in  a 
similar  way,  but  lower  temperatures  are  required  to  effect  the 
changes,  which  take  place  more  readily  in  the  case  of  calcium. 

The  solubilities  of  the  corresponding  salts  of  the  three  ele- 
ments change  progressively.  In  the  case  of  the  hydroxides  the 
solubility  decreases  in  the  order  barium,  strontium,  calcium,  the 
weight  of  the  base  dissolving  in  100  grams  of  water  at  18°  being, 
respectively,  3.7,  0.77,  and  0.17grams.  The  values  of  the  solubilities 
of  the  sulphates  are  progressive,  but  in  the  reverse  order;  they  are, 
beginning  with  barium  sulphate,  0.0s23,  0.011,  and  0.20  gram  in 
100  grams  of  water  at  18°.  The  carbonates  of  the  three  elements 

531 


532  INORGANIC  CHEMISTRY  FOR  COLLEGES 

are  exceedingly  insoluble  and  differ  but  little  from  one  another  in 
this  respect;  the  solubility  of  calcium  carbonate  is  given  as  0.0s  13 
gram  in  100  grams  of  water. 

CALCIUM 

624.  The  more  important  forms  in  which  calcium  occurs  in 
nature  have  been  stated  (556) ;  over  3  per  cent  of  the  earth's  crust 
is  calcium,  which  stands  fifth  in  the  list  of  elements  when  they  are 
arranged  according  to  their  abundance  (556) . 

Calcium  was  first  isolated  in  an  impure  condition  by  Davy  in 
1808  and  was  prepared  in  1898  by  Moissan  by  heating  calcium 
oxide  with  sodium.  It  is  manufactured  now  by  the  electrolysis  of 
fused  calcium  chloride. 

The  more  important  physical  properties  of  calcium  are  given 
in  the  table  on  page  443.  It  is  an  excellent  conductor  of  elec- 
tricity. When  heated  it  unites  with  hydrogen  to  form  calcium 
hydride,  CaEb,  and  with  nitrogen  to  form  a  nitride,  Ca3N2, 
which  decomposes  with  water  to  form  ammonia  and  calcium 
hydroxide.  The  metal,  which  has  become  only  recently  available 
commercially,  has  not  yet  found  any  industrial  uses. 

625.  Calcium  Chloride. — Sea-water  contains  calcium  chloride 
and,  consequently,  the  compound  is  found  in  the  salt  deposits 
formed  as  the  result  of  the  evaporation  of  inland  seas.     It  occurs 
in  combination  with  magnesium  chloride  in  the  mineral  tachydrite, 
CaCl2,2MgCl2,12H2O,  at  Stassfurt. 

Calcium  chloride  crystallizes  as  a  hexalrydrate,  CaCl2,  6H2O, 
from  water.  The  salt  is  very  deliquescent,  and  on  account  of  its 
great  solubility  it  is  used  in  freezing  mixtures.  When  heated  to  a 
sufficiently  high  temperature  it  loses  its  water  of  crystallization, 
but  is  at  the  same  time  partially  hydrolyzed  and  the  resulting 
product  contains  some  calcium  hydroxide  or  calcium  oxide.  In 
preparing  the  desiccated  material  to  be  used  as  a  drying  agent  the 
hexahydrate  is  heated  not  higher  than  200°  and  converted  into 
the  dihydrate,  CaCl2,  2H20,  which  is  the  so-called  granulated  cal- 
cium chloride  used  in  the  laboratory.  Even  this  contains  some 
hydroxide  and  if  it  is  to  be  used  in  the  quantitative  determination 
of  water  when  carbon  dioxide  is  present,  the  salt  is  first  exposed  to 
a  stream  of  dry  carbon  dioxide  to  convert  the  hydroxide  into  car- 


CALCIUM,  STRONTIUM,  BARIUM,  AND  RADIUM        533 

bonate.  Dehydrated  calcium  chloride  unites  with  ammonia  to 
form  a  salt,  CaChjSNHs,  with  organic  derivatives  of  water, 
such  as  the  alcohols,  and  with  the  compounds  derived  from  ammo- 
nia. For  this  reason  it  cannot  be  used  to  dry  these  substances. 

626.  Calcium  Carbonate. — Limestone,  marble,  coral,  and  chalk 
occur  in  vast  quantities  on  the  earth's  surface;  the  mountains  of 
Switzerland  are  composed  largely  of  limestone  and  those  of  the 
Dolomite  region  in  Central  Europe  of  a  mineral  which  is  called 
dolomite,  CaCOsjMgCOs,  from  the  name  of  the  scientist  who  first 
described  the  region  fully.  Limestone  in  the  purest  condition 
consists  of  minute  crystals  of  calcium  carbonate  which  are  asso- 
ciated with  varying  amounts  of  iron  oxide,  silica,  clay,  magnesium 
carbonate,  and  other  substances;  its  color  may  be  white,  gray, 
blue,  or  black  depending  upon  the  impurities  present.  Limestones 
have  been  produced  from  the  remains  of  animal  organisms  which 
existed  in  the  sea;  they  contain,  at  times,  shell  formations,  which 
can  be  clearly  seen.  The  harder  varieties  of  limestone  are  exten- 
sively used  as  a  building  material,  but  they  are  attacked  slowly 
by  air  and  water  and  are  not  as  permanent  as  sandstone  or  granite 
for  this  purpose. 

Marble  is  a  more  or  less  pure,  highly  crystalline,  massive 
variety  of  calcium  carbonate,  which  can  be  polished.  Onyx 
marbles  are  translucent  and  are  made  up  of  bands,  colored  by 
impurities,  which  have  been  deposited  from  cave  or  spring  waters. 
Marble  is  an  excellent  building  stone;  as  it  is  much  purer  than 
many  limestones  it  is  not  readily  attacked  by  atmospheric  influ- 
ences. The  red,  pink,  or  green  color  of  certain  marbles  is  largely 
due  to  iron  oxide;  the  gray  and  black  marbles  are  colored  by 
carbonaceous  material. 

Chalk  and  coral  are  made  up  of  calcium  carbonate  derived 
from  the  calcareous  parts  of  marine  organisms.  Calcium  car- 
bonate is  abundant  in  parts  of  England,  and  forms  the  well-known 
chalk  cliffs  along  its  southern  coast.  The  shells  of  eggs,  oysters, 
clams,  and  mussels  are  composed  largely  of  calcium  carbonate. 
Pearls  are  made  up  of  the  carbonate  deposited  in  exceedingly  thin 
layers  in  such  a  way  that  when  light  falls  upon  them  it  is  broken 
up  into  the  colors  of  the  spectrum. 

Calcium  carbonate  occurs  in  well-defined  individual  crystals 
as  calcite,  Iceland  spar,  and  aragonite.  The  latter  belongs  to  the 


534  INORGANIC  CHEMISTRY  FOR  COLLEGES 

rhombic  system,  and  is  an  unstable  variety;  when  heated  gently 
it  changes  to  a  mass  of  minute  crystals  of  calcite.  Iceland  spar 
and  calcite  belong  to  the  hexagonal  system.  When  calcium  car- 
bonate is  precipitated  it  is  formed  in  an  amorphous  condition,  but 
on  standing  it  becomes  crystalline  and  less  soluble  in  water.  When 
substances  are  precipitated  in  making  quantitative  analyses,  the 
precipitate  is  usually  heated  for  some  time  in  contact  with  water 
in  order  to  convert  it  into  the  crystalline  form  in  which  it  is  less 
soluble;  owing  to  its  granular  condition,  the  precipitate  is  more 
readily  separated  from  the  solution  by  filtration. 

Calcium  carbonate  is  used  in  the  manufacture  of  lime,  glass, 
and  sodium  carbonate,  as  a  flux  in  metallurgical  operations  (557), 
and  is  added  to  soils  which  have  become  acid  in  reaction  as  the 
result  of  the  formation  of  organic  acids  from  decaying  vegetable 
matter.  Calcium  carbonate  is  a  by-product  of  many  chemical 
industries.  It  is  sold  under  the  name  of  "  whiting,"  which  is 
prepared  by  grinding  and  washing  chalk,  and  is  used  to  modify  the 
shade  of  pigments,  in  polishing  powders,  and  for  other  purposes; 
mixed  with  about  18  per  cent  of  linseed  oil  it  forms  putty.  Pre- 
cipitated calcium  carbonate  is  the  chief  ingredient  of  tooth  powders. 

627.  Calcium  Bicarbonate. — Calcium  carbonate  dissolves  in 
water  containing  carbon  dioxide  as  the  result  of  the  formation  of 
an  acid  salt: 

CaC03  +  H2C03  ^  Ca(HCO3)2       • 

The  reaction  is  a  reversible  one;  the  acid  salt  is  formed  at  ordinary 
temperatures,  but  if  the  solution  containing  it  is  heated  or  evap- 
orated the  carbonate  is  precipitated  and  the  carbonic  acid  breaks 
down  into  carbon  dioxide  and  water. 

The  reaction  is  often  observed  in  regions  containing  limestone 
or  marble.  Carbon  dioxide  is  frequently  formed  within  the  earth 
and  is  found  in  the  waters  of  effervescing  springs.  When  such  water 
comes  into  contact  with  masses  of  calcium  carbonate  the  latter 
is  slowly  dissolved;  some  caves  have  been  formed  in  this  way. 
The  beautiful  formations  in  the  caves  found  in  limestone  regions 
have  been  produced  by  the  decomposition  of  the  acid  carbonate 
contained  in  the  water  which  has  dissolved  calcium  carbonate. 
As  such  water  drops  from  the  roof  of  the  cave  it  slowly  evaporates, 
and  the  deposits  of  calcium  carbonate  formed  are  called  stalactites 


CALCIUM,  STRONTIUM,  BARIUM,  AND  RADIUM        535 

if  they  hang  from  the  roof,  or  stalagmites  when  they  accumulate 
on  the  floor. 

628.  Hard  Water.  —  The  action  of  carbon  dioxide  and  water  on 
magnesium  carbonate  and  ferrous  carbonate  is  similar  to  that  just 
described  in  the  case  of  the  calcium  salt.     Water  which  contains 
salts  of  calcium,  magnesium,  or  iron  in  solution  is  said  to  be 
"  hard,"  because  when  it  is  used  with  soap  the  solution  does  not 
have  the  smooth  feeling  characteristic  of  soap  solutions.     When 
hard  water  of  this  type  is  boiled,  the  bicarbonates  present  are 
decomposed,  the  insoluble  normal  carbonates  are  precipitated, 
and  the  carbon  dioxide  escapes  as  a  gas;  the  water  is  thus  rendered 
"soft"  and  dissolves  soap.     Hard  water  which  can  be  made  soft 
by  boiling  is  said  to  be  "  temporarily"  hard. 

Calcium  sulphate  occurs  as  a  mineral  and  as  it  is  slightly 
soluble  in  water  it  is  often  found  in  water-supplies.  In  this  case 
boiling  the  water  does  not  remove  the  salt  and  for  this  reason 
the  water  is  said  to  be  "  permanently"  hard. 

Hard  water  of  either  type  cannot  be  used  with  soap  for  cleansing 
purposes  until  the  hardness  is  removed.  When  soap  dissolves  in 
water,  the  solution  produced  converts  the  grease  present  into  an 
emulsion  which  consists  of  globules  of  microscopic  size  that  stay  sus- 
pended in  the  water.  There  is  formed  what  appears  to  be  a  cloudy 
solution,  which  can  be  washed  away  by  more  water.  If  the  water 
is  hard,  the  salts  in  solution  react  with  the  soap  and  form  with  it 
insoluble  compounds  which  are  precipitated.  As  more  soap  is 
added  the  process  continues  until  the  metallic  ions  have  been 
precipitated;  the  soap  then  stays  in  solution  and  produces  its 
normal  effect  on  the  grease. 

629.  Common  soap  consists  of  a  mixture  of  the  sodium  salts 
of  certain   organic  acids  which  are  obtained  from  fats   (502); 
among  these  are  stearic  acid,  which  will  serve  as  an  example.     The 
salts  of  this  acid  which  contain  calcium,  magnesium,  or  iron  are 
insoluble  in  water,  and,  as  a  consequence,  when  soap  is  added  to 
hard  water  a  double  decomposition  takes  place,  which  is  repre- 
sented in  the  case  of  sodium  stearate  as  follows: 

CaSO4  +  2Na(Ci8H35O2)  =  Ca(Ci8H35O2)2  +  Na2SO4 


Calcium  stearate  is  precipitated  when  the  soap  is  added  to  water 
as  long  as  the  latter  contains  a  calcium  salt. 


536  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  removal  of  the  metallic  salts  from  the  water  by  means  of 
soap  takes  time  and  is  expensive,  and  other  substances  are  com- 
monly used  for  this  purpose.  If  the  water  is  temporarily  hard 
it  can  be  softened,  as  we  have  seen,  by  boiling,  but  if  large  amounts 
of  water  are  to  be  used  in  a  laundry  the  cost  of  heating  the  water  is 
prohibitive.  Water  containing  bicarbonates  can  be  softened  by 
adding  to  it  just  enough  calcium  hydroxide  to  convert  the  acid 
salts  into  neutral  salts: 

Ca(HCO3)2  +  Ca(OH)2  =  2CaCO3  +  2H20 

An  excess  must  be  avoided,  because  calcium  hydroxide  is  soluble 
and  its  presence  in  water  renders  it  hard.  Lime  is  employed  to 
soften  water  to  be  used  in  boilers.  Ammonium  hydroxide  is 
commonly  used  in  the  household;  in  this  case  the  reaction  which 
occurs  is  as  follows: 

Ca(HCO3)2  +  2NH4OH  =  CaC03  +  (NH4)2CO3  +  2H2O 

Permanently  hard  water  may  contain  in  addition  to  calcium 
sulphate  the  soluble  salts  of  other  metals;  sea- water  is  hard  because 
there  is  present  in  it,  along  with  other  substances,  the  chlorides  of 
calcium  and  magnesium.  In  order  to  soften  hard  water  of  this 
type,  it  is  treated  with  the  sodium  salt  of  an  acid  which  forms 
insoluble  calcium  and  magnesium  salts.  A  carbonate  or  borate 
(borax)  can  be  used  for  this  purpose: 

CaSO4  +  Na2CO3  =  CaCO3+  Na2SO4 

630.  In  the  "  permutit  "  process  for  softening  water,  the  water 
is  caused  to  flow  slowly  over  an  artificial  silicate  which  approaches 
in  composition  the  formula  NaAlSi04,3H2O.  Silicates  of  similar 
composition,  which  are  derivatives  of  orthosilicic  acid,  H4SiO4, 
occur  in  nature  and  are  called  zeolites.  They  are  much  more 
readily  attacked  by  water  than  feldspar,  KAlSi3Og,  and  similar 
minerals  which  are  richer  in  silica.  Permutit  is  furnished  in  the 
form  of  flat  granules  about  one-quarter  of  an  inch  in  diameter. 
When  it  comes  into  contact  with  hard  water  a  double  decomposi- 
tion takes  place,  sodium  passes  into  solution,  and  is  replaced  in  the 
silicate  by  calcium  and  magnesium.  When  the  permutit  no 
longer  acts  in  this  way,  it  is  treated  with  a  strong  solution  of  salt; 
the  reaction  is  reversed  and  the  material  regenerated.  The 


CALCIUM,  STRONTIUM,  BARIUM,  AND  RADIUM        537 

process  is  a  striking  practical  example  of  the  application  of  the  law 
of  molecular  concentration  to  reversible  reactions. 

631.  The  results  of  the  analysis  of  a  hard  water  are  usually 
expressed  in  parts  per  million,  the  substances  producing  hardness 
being  calculated  as  calcium  carbonate.     It  is  also  expressed  in 
"  degrees,"  which  in  the  United  States  indicate  the  number  of 
grains  per  U.  S.  gallon:  1  degree  equals  17.1  parts  per  million. 
Hard  water  from  wells  in  limestone  regions  may  contain  as  much 
as  400  parts  per  million.     The  action  of  hard  water  in  boilers  will 
be  considered  later. 

632.  Calcium   Oxide. — Quicklime,   or  lime,   CaO,   is   a    very 
important  industrial  product  and  is  made  in  large  quantities  by 
heating  limestone: 

CaC03  ^±  CaO  +  C02 

The  reaction  is  a  reversible  one  and  the  equilibrium  attained  is 
determined  by  the  temperature,  provided  the  gas  does  not  escape. 
At  700°  the  pressure  of  carbon  dioxide  in  equilibrium  with  the  oxide 
is  25  mm.  and  at  900°  it  reaches  1  atmosphere.  In  making  lime, 
calcium  carbonate  is  "  burned  "  at  about  1000°  in  kilns  40  to  50 
feet  high  and  6  to  10  feet  in  diameter,  constructed  of  brick  or  of 
iron  plates  lined  with  fire-brick.  The  limestone  is  added  at  the 
top  of  the  kiln  and  the  lime  withdrawn  from  the  bottom.  Since 
lime  and  calcium  carbonate  are  poor  conductors  of  heat  the  fuel 
selected  for  use  is  that  which  produces  a  flame  which  penetrates 
some  distance  into  the  kiln;  wood,  soft  coal,  oil,  or  gas  are  used  for 
this  purpose.  An  ingenious  device  is  now  used  in  certain  kilns 
which  markedly  simplifies  the  lime-burning.  The  gases  are  drawn 
off  from  the  top  of  the  kiln  by  means  of  a  fan  and  a  part  of  them  is 
forced  into  the  air  supply  furnished  the  grate  where  the  fuel  is 
burning.  Dilution  of  the  air  with  carbon  dioxide  lowers  the  tem- 
perature of  the  flame  (230)  and  increases  its  size;  as  the  result  of 
the  "  long-flame  "  combustion  of  the  fuel  a  more  even  heating  of 
the  charge  at  a  lower  temperature  is  secured  (233). 

The  limestone  used  in  making  lime  is  more  or  less  contam- 
inated with  magnesium  carbonate,  silica,  clay,  oxides  of  iron,  and 
other  substances,  and  if  it  is  heated  during  its  formation  above 
1200°,  the  calcium  oxide  and  the  silica  react;  fusible  silicates 
are  formed  which  fill  up  to  some  extent  the  pores  in  the  calcium 


538  INORGANIC  CHEMISTRY  FOR  COLLEGES 

oxide.  Such  lime  is  said  to  be  "  over-burned  "  and  reacts  slowly 
with  water. 

A  lime  which  contains  less  than  5  per  cent  of  impurities  is 
called  "  fat  "  or  "  rich."  Magnesium  limes  may  contain  as  much 
as  30  per  cent  of  magnesium  oxide;  they  react  with  water  slowly, 
and  yield  a  mixture  of  calcium  and  magnesium  hydroxides,  which 
produce  a  " smooth  "  texture;  limes  of  this  kind  are  used  in  plasters 
for  finishing  purposes. 

633.  Slaked  Lime. — The  addition  of  water  to  calcium  oxide  to 
form  calcium  hydroxide  is  called  "  slaking  ": 

CaO  +  H2O  =  Ca(OH)2 

The  union  is  accompanied  by  the  evolution  of  heat,  which  in  con- 
fined places  may  result  in  fire. 

In  slaking  lime  to  prepare  mortar  the  former  is  treated  with 
about  twice  the  amount  of  water  required  to  convert  it  into  cal- 
cium hydroxide.  The  thick  paste  formed  when  the  lime  has 
slaked  is  mixed  with  2.5  to  3  volumes  of  sand,  and  water  is  added 
to  give  the  mortar  the  right  consistency. 

When  mortar  sets  it  first  loses  the  water  it  contains  and  then 
slowly  absorbs  carbon  dioxide  from  the  air;  the  calcium  carbonate 
formed  crystallizes  around  the  grains  of  sand  and  binds  them 
together.  Calcium  carbonate  is  formed  to  a  depth  of  about  0.1 
inch  during  the  first  year,  but  as  it  increases  in  thickness  the 
change  takes  place  more  slowly.  About  twenty-three  years  are 
required  to  convert  the  mortar  in  an  ordinary  brick  wall  into  cal- 
cium carbonate.  The  presence  of  sand  in  mortar  renders  the 
mixture  more  or  less  porous  when  it  dries,  and  the  absorption 
of  carbon  dioxide  is,  therefore,  facilitated. 

Lime  that  is  exposed  to  the  air  absorbs  carbon  dioxide,  and 
when  it  is  converted  in  this  way  into  calcium  carbonate  it  will  not 
react  with  water;  it  is  said  to  be  "  air-slaked." 

634.  Calcium  Hydroxide. — The  way  in  which  calcium  hydrox- 
ide is  prepared  from  lime  has  just  been  described  (633).     The 
compound  finds  many  important  uses  in  industrial  chemistry 
because  it  is  the  cheapest  base.     It  dissolves  slightly  in  water  (1 
part  in  600  parts  at  18°);   the  solution,  called  limewater,  is  used 
in  medicine  when  a  very  dilute  solution  of  an  alkali  is  needed.    A 


CALCIUM,  STRONTIUM,  BARIUM,  AND  RADIUM        539 

suspension  of  the  finely  divided  solid  in  water  is  known  as  "  milk 
of  lime." 

Calcium  hydroxide  is  used  in  making  the  caustic  alkalies  and 
bleaching  powder;  for  purifying  illuminating  gas;  in  the  manu- 
facture of  many  chemicals,  such  as  acetic  acid;  in  purifying  sugar 
solutions;  in  bleaching  cotton;  in  tanning  to  remove  the  hair 
from  hides,  etc. 

635.  Calcium  Sulphate. — A  hydrated  form  of  calcium  sulphate, 
CaSO4,2H2O,   occurs  as  the  mineral  gypsum.     It  is  mined  in 
large  quantities  and  is  used  in  making  plaster  and  stucco.     A  pure 
granular  form  of  the  dihydrate  is  called  alabaster,  and  as  it  takes  a 
high  polish  it  is  used  in  making  statuary.     The  solubility  of  gypsum 
in  water  is  0.2  gram  in  100  grams. 

When  gypsum  is  heated  it  loses  water  and  is  first  converted 
into  a  hemihydrate — 

2(CaSO4,2H2O)  ^  (CaSO4)2,H2O  -f-  3H2O 

three-quarters  of  the  water  being  lost.  At  about  200°  the  rest  of 
the  water  is  given  off  and  the  anhydrous  form  of  the  salt  is  obtained. 
The  hemihydrate  when  mixed  with  water  unites  with  it  readily, 
whereas  the  union  of  anhydrous  calcium  sulphate  and  water  takes 
place  very  slowly. 

636.  Plaster  of  Paris,  which  is  the  hemihydrate,  (CaSO4)2,H2O, 
is  manufactured  in  large  quantities  by  calcining  gypsum  at  about 
145°  in  muffles  or  rotary  cylinders.     When  the  finely  ground 
material  is  mixed  with  about  twice  the  theoretical  quantity  of 
water  required  to  convert  it  into  gypsum,  it  forms  a  plastic  mass 
which  "sets"  in  a  short  time  to  a  solid.     Since  the  mass  expands 
on  setting,  plaster  of  Paris  can  be  cast  in  molds,  and  for  this 
reason  it  is  used  in  making  casts  of  statuary,  wall  decorations, 
etc.     These  are  frequently  treated  with  a  solution  of  paraffin 
or  stearin  in  gasoline,  which  leaves  on  evaporation  a  thin  coating 
of  the  wax  that  fills  the  pores  of  the  plaster.     This  treatment 
prevents  the  solution  of  the  plaster  by  water  and  the  accumulation 
of  dust.     Solutions  of  this  kind  are  also  valuable  for  water-proofing 
articles  made  of  textile  fabrics,  such  as  tents,  etc. 

The  setting  of  plaster  of  Paris  results  from  the  fact  that  when 
it  is  mixed  with  water  the  latter  dissolves  some  of  the  hemihydrate 
which  unites  with  water  to  form  the  dihydrate.  Since  the  latter 


540  INORGANIC  CHEMISTRY  FOR  COLLEGES 

is  less  soluble  than  the  former  it  crystallizes  from  the  solution, 
and  more  of  the  hemihydrate  dissolves;  the  process  continues 
until  the  plaster  has  been  changed  into  a  mass  of  interlocking 
crystals  from  which  the  excess  of  water  evaporates.  If  the 
plaster  has  been  "hard-burned"  or  " dead-burned "  it  requires 
hours  or  days  for  the  setting,  which  takes  place  in  the  case  of  ordi- 
nary plaster  of  Paris  in  a  few  minutes;  the  product  obtained  from 
dead-burned  plaster  is  hard  and  strong  and  is  preferred  for  floors 
and  for  hard-finish  plasters. 

The  fact  that  the  presence  of  colloidal  substances  retards 
crystallization  is  utilized  when  it  is  advisable  to  have  the  setting 
of  plaster  of  Paris  take  place  slowly.  This  is  accomplished  by 
adding  to  the  plaster,  glue,  vegetable  gums,  or  fine  saw-dust; 
such  a  mixture  is  called  "  stucco."  The  name  stucco  is  also 
applied  to  the  mortar  finish  put  on  the  exterior  walls  of  buildings; 
it  is  a  cement-sand  mortar  containing  some  fiber  to  give  it 
strength.  Wall  plaster  contains  in  addition  to  the  retarding  agent 
about  2  pounds  of  hair  per  ton,  which  give  additional  strength  to 
the  plaster. 

The  precipitated  calcium  sulphate  which  is  a  by-product  hi  a 
number  of  chemical  industries,  is  used  under  the  names  "  crown 
filler  "  and  "  pearl  hardening,"  as  a  filler  in  paper  making  and  in 
weighting  cloth. 

637.  The  Phosphates  of  Calcium. — The  tertiary  salt  of  ortho- 
phosphoric  acid,  Cas(P04)2,  occurs  widely  distributed  as  phos- 
phorite and,  in  combination  with  calcium  fluoride  or  chloride,  as 
apatite,  CaF2(Cl2),3Ca3(PO4)2.  The  ashes  derived  from  bones 
contain  about  80  per  cent  of  calcium  phosphate.  The  salt  is 
precipitated  when  a  soluble  phosphate  is  added  to  a  solution  of  a 
calcium  salt;  it  is  soluble  in  nitric  acid  (593),  and  is  reprecipitated 
when  the  solution  is  neutralized.  It  is  converted  by  sulphuric 
acid  into  calcium  sulphate  and  the  secondary  salt,  CaHPO4,  the 
primary  salt,  Ca(H2PO4)2,  or  phosphoric  acid,  H3PO4,  the  extent 
of  the  conversion  being  determined  by  the  relative  amounts  of  the 
two  substances  used.  The  secondary  salt  is  difficultly  soluble  in 
water  (0.23  gram  per  liter);  the  primary  salt  is  more  soluble  (18 
grams  per  liter).  Primary  calcium  phosphate  is  made  in  the  pure 
condition  for  use  in  baking  powders,  and  the  mixture  of  the  salt 
with  calcium  sulphate  and  some  secondary  phosphate  is  manu- 


CALCIUM,  STRONTIUM,  BARIUM,  AND  RADIUM        541 

factured  on  the  large  scale  for  use  as  a  fertilizer  and  sold  under  the 
name  "  super-phosphate." 

638.  Fertilizers. — The  importance  of  ammonia  or  nitrates  and 
of  potassium  salts  in  the  growth  of  plants  has  been  emphasized. 
In  addition  to  these  substances  a  growing  plant  takes  from  the  soil 
phosphates,  silicon,  iron,  sodium,  magnesium,  chlorine,  and  other 
elements.  Most  soils  contain  an  adequate  supply  of  all  these 
except  phosphates,  which  are  present  in  certain  regions  in  limited 
amounts  only.  When  any  of  the  plant-foods  become  deficient 
they  must  be  added  to  the  soil  if  it  is  to  retain  its  fertility.  Nitro- 
gen is  supplied  as  sodium  nitrate,  ammonium  sulphate,  calcium 
nitrate,  calcium  cyanamide  (342),  or  in  the  form  of  guano,  manure 
or  sewage;  the  source  of  potassium  is  the  chloride  or  sulphate,  and 
that  of  phosphoric  acid  one  or  more  of  the  calcium  phosphates. 
A  "  complete  "fertilizer  contains  substances  which  furnish  nitro- 
gen, potassium,  and  phosphorus,  but  as  many  soils  lack  only  one 
of  the  plant  foods,  the  materials  are  used  separately  when  the 
nature  of  the  soil  is  known.  The  soil  in  regions  that  contain 
granitic  rocks  are  usually  rich  in  potassium,  which  is  derived  from 
feldspar;  other  soils  are  supplied  with  phosphates  from  the  phos- 
phorite they  contain. 

Lime  or  finely  powdered  calcium  carbonate  is  added  to  soils 
which  have  become  too  acid  as  the  result  of  the  disintegration  of  the 
organic  matter  present,  which  forms  the  so-called  " humus." 
Calcium  sulphate  is  also  added  to  acid  soils.  It  is  believed  that 
the  salt  serves  to  fix  the  ammonia  contained  in  the  air.  Ammonia 
and  carbon  dioxide  form  ammonium  carbonate,  which  inter- 
acts with  the  sulphate  to  produce  calcium  carbonate  and  ammo- 
nium sulphate.  The  former  neutralizes  the  organic  acids,  and  the 
latter,  being  more  stable  than  ammonium  carbonate,  remains  in 
the  soil  and  thus  furnishes  a  supply  of  fixed  nitrogen. 

Garbage  is  now  treated  in  large  cities  to  recover  the  valuable 
products  it  contains.  It  is  heated  with  steam  under  pressure 
and  the  soft  pulp  produced  is  pressed ;  the  oil  is  separated  and  used 
for  soap  and  candle  stock,  and  the  press  cake  is  dried  and  used  as 
a  "  filler  "  in  fertilizers,  since  it  contains  nitrogen,  phosphorus, 
and  a  little  potassium. 

Bones  which  have  been  heated  with  steam  to  remove  organic 
material  are  ground  and  used  as  fertilizers,  but  as  the  solubility 


542  INORGANIC  CHEMISTRY  FOR  COLLEGES 

of  tricalcium  phosphate  is  very  small,  bone-meal  is  not  a  good 
fertilizer.  The  organic  matter  contained  in  bones  interferes 
with  its  solution  in  the  soil,  and  it  is  at  times  removed  by  heating 
the  bones  to  a  high  temperature  in  the  presence  of  air.  To  render 
the  material  more  soluble,  it  is  usually  converted  into  super- 
phosphate with  sulphuric  acid. 

The  most  important  sources  of  phosphoric  acid  for  fertilizers 
are  the  great  beds  of  phosphorite  which  are  more  or  less  widely 
distributed.  Those  occurring  in  Florida,  South  Carolina,  and  Ten- 
nessee are  placed  so  they  are  easy  to  mine,  and  because  they  are 
porous  they  are  readily  handled. 

The  ground  rock  is  treated  with  enough  "  chamber  "  acid  to 
react  according  to  the  following  equation: 

Ca3(PO4)2  +  2H2S04  +  6H2O  = 

Ca(H2PO4)2,2H2O  +  2(CaSO4,2H2O) 

The  mixture  produced  contains  some  phosphoric  acid  in  addition 
to  the  primary  salt;  it  is  sold  under  the  name  of  superphosphate. 
If  the  conversion  of  the  tertiary  salt  has  not  been  completely 
effected,  the  part  left  in  the  mixture  reacts  on  standing  with  some 
of  the  primary  salt  to  form  the  secondary  salt: 

Ca3(PO4)2  +  Ca(H2PO4)2  =  4CaHPO4 

A  similar  reaction  takes  place  between  the  tertiary  salt  and  the 
free  phosphoric  acid : 

Ca3(PO4)2  +  H3PO4  =  3CaHPO4 

These  reactions  constitute  what  is  called  "  reversion."  Since 
secondary  calcium  phosphate  is  very  difficultly  soluble  in  water,  a 
"  reverted  "  phosphate  is  absorbed  much  more  slowly  by  the  soil 
and  is  much  less  valuable  as  a  fertilizer. 

639.  Bleaching  Powder. — This  compound,  which  is  also  called 
chloride  of  lime  or  chlorinated  lime,  is  manufactured  by  treating 
slaked  lime  with  chlorine,  either  in  absorption  chambers  built  of 
brick,  cast-iron,  or  lead,  or  in  cast-iron  cylinders  provided  with 
rotating  blades  which  act  as  conveyors  and  keep  the  solid  in  motion 
so  that  it  comes  in  contact  with  the  gas.  The  chlorine  must  be 
admitted  slowly  as  the  reaction  which  takes  place  is  exothermic 
and  a  rise  in  temperature  interferes  with  the  absorption.  The 


CALCIUM,  STRONTIUM,  BARIUM,  AND  RADIUM        543 

materials  are  left  in  contact  until  the  bleaching  powder  formed 
contains  from  36  to  37  per  cent  "  available  "  chlorine,  that  is, 
chlorine  which  is  liberated  when  the  powder  is  treated  with  an  acid. 
Bleaching  powder  is  a  mixed  salt  of  hydrochloric  and  hypochlorous 

,c\ 

acids  and  has  the  formula  Ca<T         .     It  is  not  a  mixture  of  calcium 

XOC1 

chloride  and  calcium  hypochlorite,  because  it  does  not  show  the 
properties  characteristic  of  calcium  chloride,  which  deliquesces 
and  is  soluble  in  alcohol. 

Bleaching  powder  absorbs  carbon  dioxide  and  water  from  the 
air,  and  the  carbonic  acid  formed  liberates  hypochlorous  acid, 
which  produces  the  characteristic  odor  of  the  powder.  It  is  used 
extensively  in  bleaching  vegetable  fibers  in  the  paper  and  textile 
industries.  The  material  is  first  passed  through  a  solution  of  the 
salt  and  then  into  a  dilute  solution  of  an  acid  to  set  free  chlorine, 
which  bleaches  as  the  result  of  the  formation  of  hypochlorous  acid. 
Bleaching  powder  is  also  used  as  a  disinfectant,  as  it  destroys  the 
bacteria  which  bring  about  putrefaction. 

640.  Calcium  Sulphide. — When  calcium  sulphate  is  heated  with 
carbon  to  a  high  temperature  it  is  reduced  to  calcium  sulphide: 

CaSO4  +  4C  =  CaS  +  4CO 

The  salt  is  difficultly  soluble  in  water  but  is  slowly  hydrolyzed: 
2CaS  +  2H2O  ^±  Ca(OH)2  +  Ca(SH)2 

The  hydrosulphide  and  a  part  of  the  hydroxide  pass  into  solution. 
The  sulphide  is  used  to  dissolve  hair,  and  in  the  preparation  of 
luminous  paints.  When  exposed  to  sunlight  calcium  sulphide 
acquires  the  property  of  giving  off  light  in  the  dark.  The  cause 
of  the  phenomenon  is  not  known,  but  is,  in  all  probability,  due 
to  the  presence  of  traces  of  impurities.  A  solution  of  calcium 
polysulphide  is  used  as  a  tree-spray  to  destroy  fungi  and  insects. 

641.  Glass. — Calcium  silicate  is  a  constituent  of  the  variety  of 
glass  used  to  make  bottles  and  window  panes.     Glass  of  this  kind  is 
prepared  by  heating  together  to  a  high  temperature  limestone, 
sand,  and  sodium  carbonate  or  sodium  sulphate,  until  a  homo- 
geneous, transparent  liquid  mass  is  obtained.     The  product  is 
then  blown  into  the  shape  desired;  if  bottles  are  to  be  made  some 


544  INORGANIC  CHEMISTRY  FOR  COLLEGES 

of  the  molten  glass  adhering  to  the  end  of  a  hollow  tube  is  placed 
in  a  mold  and  air  is  blown  through  the  tube  until  the  glass  assumes 
the  form  of  the  mold.  If  window  glass  is  to  be  made,  the  glass  is 
blown  into  the  form  of  a  long  cylinder;  when  it  has  cooled  the 
ends  of  the  latter  are  cut  off,  the  glass  is  cut  in  a  line  parallel  to  the 
length  of  the  cylinder,  and  after  heating  it  until  it  softens  the 
cylinder  is  opened  up  to  form  a  flat  piece  of  glass.  Plate  glass  is 
made  by  rolling  out  the  hot  material  into  the  form  of  a  sheet,  and 
grinding  and  polishing  the  surface  of  the  glass  while  it  moves  on  a 
rotating  table. 

Glass  for  these  purposes  has  approximately  the  composition 
represented  by  the  formula  Na2O,CaO,6SK)2.  It  is  a  mixture  of 
silicates  of  sodium  and  calcium  which  contain  a  higher  percentage 
of  silicon  dioxide  than  is  present  in  the  normal  silicates  of  these 
metals.  Glass  is  a  supercooled  liquid,  because  it  is  cooled  so 
rapidly  from  the  molten  condition  that  the  transition  point  at 
which  the  silicates  change  from  the  liquid  to  the  solid  crystalline 
condition  is  passed  before  the  change  can  take  place.  The  result- 
ing glass  is  so  hard  that  it  remains  in  this  condition  for  a  very  long 
time  because  the  materials  are  prevented  from  taking  up  the 
orderly  arrangement  that  exists  in  a  crystal.  Some  varieties  of 
glass  on  standing  become  more  or  less  opaque  as  a  result  of  the 
partial  crystallization  of  its  constituents. 

Glasses  which  possess  widely  different  physical  and  chemical 
properties  can  be  made  by  substituting  in  part  or  in  whole,  sodium 
and  calcium  by  other  metallic  elements,  and  silicon  dioxide  by 
other  acid  anhydrides.  If  sodium  is  replaced  by  potassium  a  glass 
is  obtained  which  melts  at  a  much  higher  temperature  than  bottle 
glass.  If  lead  replaces  calcium  the  glass  has  a  high  index  of  refrac- 
tion which  makes  it  of  particular  value  in  the  construction  of 
lenses  for  optical  purposes.  Such  glasses  may  contain  also  barium 
or  zinc. 

Ordinary  glass  is  slowly  attacked  by  hot  water  and  acids  as  the 
result  of  the  decomposition  of  the  silicates  which  it  contains.  For 
this  reason  it  is  not  well  adapted  to  the  uses  to  which  glass  appa- 
ratus is  put  in  a  chemical  laboratory.  The  defect  is  overcome  by 
using  for  this  purpose  a  glass  containing  a  low  percentage  of 
metallic  oxides,  and,  in  addition  to  silicon  dioxide,  borrc  anhy- 
dride and  aluminium  oxide;  the  latter  oxide  renders  the  glass 


CALCIUM,  STRONTIUM,  BARIUM,  AND  RADIUM        545 

tougher,  so  that  it  is  less  liable  to  break  under  a  mechanical  strain. 
The  cubical  expansion  of  glass  made  in  this  way  is  much  less  than 
that  of  ordinary  glass  on  account  of  the  high  percentage  of  silica 
present  (507),  and,  as  a  consequence,  the  glass  does  not  crack  so 
readily  when  subjected  to  sudden  changes  in  temperature.  One 
kind  of  glass  of  this  type  now  used  has  approximately  the  following 
percentage  composition:  Na2O,  5;  SiCb,  80;  I^Os,  12,  and 
A^Os,  3.  It  is  made  by  fusing  together  borax  and  sand  which 
contains  a  small  percentage  of  aluminium  oxide. 

Glass  is  colored  by  adding  to  it  substances  which  either  form 
colored  silicates  or  dissolve  in  the  molten  glass  in  the  elementary 
condition.  Cobalt  silicate  is  blue,  ferrous  silicate  is  green,  etc. 
Selenium  furnishes  a  red  glass  and  metallic  gold  in  the  colloidal 
condition  gives  a  ruby  glass.  The  color  of  window  glass  when 
viewed  in  a  thick  layer  is  green  on  account  of  the  ferrous  silicate 
present.  If  the  material  from  which  colorless  glass  is  to  be 
made  contains  much  iron,  manganese  dioxide  is  added  to  the 
melted  glass  to  oxidize  the  iron  to  the  ferric  condition ;  the  latter  is 
light  yellow,  and  the  glass  has,  as  a  consequence,  less  color. 

642.  Tests  for  Calcium  Salts. — The  presence  of  these  salts  is 
most  readily  recognized  by  the  red  color  they  impart  to  the  Bun- 
sen  flame.  The  presence  of  calcium  ions  is  shown  by  the  precip- 
itation of  colorless  calcium  oxalate  when  the  solution  under  exam- 
ination, after  the  addition  of  acetic  acid,  is  treated  with  ammonium 
oxalate,  (NH4)2C2O4.  The  oxalates  of  strontium  and  barium 
are  insoluble  in  water,  but  are  not  precipitated  if  acetic  acid  is 
present.  In  the  quantitative  determination  of  calcium,  the  oxa- 
late is  separated,  ignited,  and  weighed  as  calcium  oxide : 

CaC2O4  =  CaO  +  CO  +  CO2 


STRONTIUM 

643.  Strontium  occurs  as  strontianite,  SrCOs,  and  as  celestite, 
SrSO4.  Its  compounds  resemble  closely  those  of  barium  (646) 
and  are  prepared  in  a  similar  way;  the  chloride,  SrCl2,6H2O,  and 
the  nitrate,  Sr(N03)2,  and  the  hydroxide,  Sr(OH)2,8H2O,  are 
used  as  laboratory  reagents.  The  nitrate  and  chlorate  are  used 
in  "  red-fires." 


546  INORGANIC  CHEMISTRY  FOR  COLLEGES 

644.  Test  for  Strontium  Salts. — Strontium  forms  an  insoluble 
carbonate,  sulphate,  and  oxalate.     Its  compounds  give  a  charac- 
teristic carmine-red  color  to  the  Bunsen  flame.     The  fact  that 
strontium  salts  are  precipitated  as  the  sulphate  by  a  solution  of 
calcium  sulphate  serves  to  identify  the  element  in  the  presence 
of  calcium.     Strontium  salts  are  not  affected  by  a  solution  of 
strontium  sulphate,  which,  however,  produces  a  precipitate  with 
barium  salts. 

BARIUM 

645.  Barium  occurs  as  witherite,  BaCOa,  and  barite  or  heavy 
spar,   BaSO4,  which  has  an  unusually  high  density,  4.5.     The 
metal  is  prepared  by  the  electrolysis  of  the  molten  chloride,  and 
its  salts  are  made  by  the  action  of  acids  on  the  carbonate  or  the 
sulphide. 

646.  The  Oxides  of  Barium. — Barium  oxide,  BaO,  is  prepared 
by  heating  the  carbonate;  as  the  temperature  is  high,  about  1500°, 
charcoal  is  added  to  help  in  the  reduction: 

BaCO3  +  C  =  BaO  +  2CO 

The  oxide  reacts  with  water  vigorously  and  is  converted  into  the 
hydroxide.  When  heated  in  a  stream  of  air  at  about  500°  it 
unites  with  oxygen  to  form  barium  peroxide,  which  is  used  in  the 
manufacture  of  hydrogen  peroxide.  A  hydrated  form  of  the 
peroxide,  BaO2,8H2O,  is  obtained  as  a  precipitate  when  hydrogen 
peroxide  is  added  to  a  solution  of  barium  hydroxide.  Calcium  and 
strontium  yield  similar  compounds,  but  the  anhydrous  peroxides 
are  not  formed  when  the  oxides  are  heated  in  air  or  oxygen. 

647.  Barium   Hydroxide. — This    compound    is   manufactured 
from  barium  oxide  and  water.     The  oxide  is  prepared  as  described 
above  or  is  obtained  by  reducing  barium  sulphate  in  an  electric 
furnace: 

BaSO4  +  4C  =  BaS  +  4CO 

3BaSO4  +  BaS  =  4BaO  +  4SO2 

The  oxide  is  formed  in  a  single  operation  and  is  tapped  from  the 
furnace  in  the  molten  condition.  It  is  dissolved  in  water  and  con- 
verted into  barium  hydroxide,  which  is  obtained  as  crystals  having 


CALCIUM,  STRONTIUM,  BARIUM,  AND  RADIUM        547 

the  composition  Ba(OH)2,81i2O.  Barium  hydroxide  is  the  most 
soluble  of  the  hydroxides  of  the  alkaline  earth  metals;  an  aqueous 
solution  known  as  "  baryta-water"  can  contain  as  much  as  3.7 
grams  of  Ba(OH)2  in  100  grams  of  water.  A  solution  of  the 
hydroxide  is  much  used  in  quantitative  volumetric  analysis. 

648.  Barium  Chloride. — This  salt  is  manufactured  from  the 
oxide  or  carbonate  and  hydrochloric  acid,  or  from  the  sulphate, 
carbon,  and  calcium  chloride: 

BaS04  +  4C  +  CaCl2  =  BaCl2  +  CaS  +  4CO 

It  differs  from  calcium  chloride  in  that  it  crystallizes  with  2  mole- 
cules of  water,  BaCl2,2H20,  and  is  stable  in  the  air;  it  can  be 
completely  dehydrated  without  being  hydrolyzed — a  fact  which 
shows  that  barium  is  a  stronger  base-forming  element  than  calcium. 

649.  Barium  Sulphate. — Barite,  BaSO4,  is  the  source  of  most  of 
the  compounds  of  barium.     In  a  finely  ground  condition  it  is  used 
as  a  pigment  when  mixed  with  white  lead.     It  has  little  covering 
capacity  and  is  to  be  considered  as  an  adulterant.     The  precip- 
itated sulphate,  obtained  as  a  by-product  in  some  chemical  indus- 
tries, is  also  used  as  a  pigment  and  is  better  for  this  purpose  than 
the  ground  mineral. 

Lithopone  is  a  white  pigment  which  is  a  mixture  of  precipitated 
barium  sulphate  and  zinc  sulphide  formed  by  heating  zinc  sulphate 
and  barium  sulphide  with  water : 

BaS  +  ZnSO4  =  ZnS  +  BaS04 

The  precipitate  is  washed,  dried,  mixed  with  a  little  ammonium 
chloride,  heated  red  hot,  and  quenched  in  water.  This  treatment 
serves  to  put  the  material  in  such  a  physical  condition  that  it  has  a 
high  covering  power,  mixes  well  with  oil,  and  makes  a  good  paint. 
Barium  sulphate,  like  the  sulphates  of  calcium  and  strontium, 
dissolves  in  concentrated  sulphuric  acid  to  form  an  acid  salt,  which 
is  converted  into  the  neutral  sulphate  when  the  solution  is  added 
to  water. 

650.  Other  Salts  of  Barium.— Barium  nitrate,  Ba(NO3)2,  is 
an  anhydrous  salt  which  is  a  laboratory  reagent.     Barium  chlorate, 
Ba(C103)2,  mixed  with  sulphur  and  charcoal  is  used  in  the  prep- 
aration of  fireworks  which  burn  with  a  green  flame.     Barium 
sulphide,  BaS,  is  used  in  luminous  paints. 


548  INORGANIC  CHEMISTRY  FOR  COLLEGES 

651.  Tests  for  Barium  Salts. — Since  barium  sulphate  is  the 
most  insoluble  of  the  sulphates  of  the  alkaline  earths,  it  is  pre- 
cipitated when  a  solution  of  strontium  sulphate  is  added  to  a 
barium  salt;   the  reaction  serves  to  distinguish  barium  from  the 
other  elements  which  yield  insoluble  sulphates.     Barium  chromate, 
BaCrC>4,  prepared  by  adding  a  soluble  chromate  to  a  barium  salt,  is 
insoluble  in  acetic  acid,  whereas  the  chromates  of  calcium  and 
strontium,  which  are  insoluble  in  water,  dissolve  in  the  acid. 

Barium  is  determined  quantitatively  by  precipitating  and 
weighing  it  as  barium  sulphate. 

RADIUM 

652.  The  remarkable  properties  of  radium  salts,  the  study  of 
which  has  had  such  an  influence  on  the  development  of  our  knowl- 
edge of  the  constitution  of  matter,  will  be  discussed  in  some  detail 
later;  at  this  point  attention  will  be  drawn  only  to  the  properties 
of  the  metal  and  its  compounds  which  resemble  those  of  the  other 
members  of  the  calcium  family. 

In  1896  Becquerel,  a  French  scientist,  discovered  that  minerals 
of  uranium  gave  off  a  radiation  that  affected  a  photographic  plate 
and  passed  through  bodies  opaque  to  light.  In  1903,  M.  and  Mme. 
Curie  showed  that  the  radiation  produced  by  minerals  containing 
uranium  was  stronger  than  that  observed  in  the  case  of  pure  salts 
of  the  metal.  It  was  concluded,  therefore,  that  the  minerals 
contained  another  substance  more  radio-active  than  uranium,  and, 
as  a  result  of  a  careful  study  of  the  problem,  a  new  element, 
which  was  named  radium,  was  discovered.  The  element  was 
separated  from  the  mineral  pitch-blende,  which  contains  uranium 
oxide,  U2O3,  along  with  a  large  number  of  compounds  of  other 
elements. 

When  the  constituents  of  the  mineral  were  separated  by  the 
methods  used  in  qualitative  analysis,  it  was  found  that  the 
barium  sulphate  obtained  was  the  most  active  of  all.  This  was 
converted  into  the  oxide  from  which  the  bromide  was  prepared. 
Repeated  fractional  crystallization  of  this  salt  served  to  separate 
it  into  barium  bromide  and  the  bromide  of  the  new  element. 
In  1910  the  metal  was  obtained  in  the  free  condition  by  elec- 
trolyzing  a  solution  of  radium  chloride,  using  mercury  as  the 


CALCIUM,  STRONTIUM,  BARIUM,  AND  RADIUM        549 

cathode.  The  solution  of  the  metal  prepared  in  this  way  was 
heated  to  remove  the  mercury  by  distillation.  The  radium 
obtained  was  found  to  be  a  white  metal,  which  melted  at  700°, 
tarnished  in  the  air,  and  decomposed  water  with  the  formation 
of  the  hydroxide  and  the  evolution  of  hydrogen. 

Radium  salts  are  now  manufactured  in  the  United  States  under 
the  supervision  of  the  Bureau  of  Mines  from  carnotite,  a  uranium 
ore  which  is  found  in  Colorado.  It  is  estimated  that  a  ton  of  the 
minerals  used  as  a  source  of  radium  contains  about  0.2  gram  of  the 
element;  it  is  for  this  reason  that  the  cost  of  pure  radium  salts 
is  so  high — about  $120,000  per  gram.  Highly  impure  salts  of 
radium  can  be  used  for  most  purposes  to  which  it  is  put,  and  the 
cost  of  the  material  is  correspondingly  less. 

Many  rocks  and  the  waters  of  some  mineral  springs  show  the 
property  of  radio-activity  to  a  slight  degree.  The  spectrum  of 
radium  consists  of  two  red  bands,  a  blue-green  band,  and  two  faint 
violet  lines.  The  properties  of  the  salts  of  the  metal,  as  far  as  they 
have  been  studied,  are  in  accord  with  the  position  of  the  element 
in  the  periodic  classification. 

EXERCISES 

1.  Summarize  in  the  form  of  a  table  the  physical  and  chemical  properties 
of  the  alkaline  earth  metals  and  their  oxides,  chlorides,  nitrates,  carbonates, 
and  sulphates.     State  what  you  think  would  be  the  properties  of  the  analo- 
gous compounds  of  radium,  taking  into  consideration  the  properties  of  the  com- 
pounds of  the  other  metals  of  the  group  and  the  position  of  radium  in  the 
family. 

2.  (a)  Why  is  a  very  soluble  salt  more  efficient  in  making  freezing  mix- 
tures than  a  less  soluble  one?     (6)  Should  CaCl2,  6H2O  or  CaCl2  be  used 
in  making  a  freezing  mixture?     Give  a  reason  for  your  answer  . 

3.  How  could  you  show  that  dolomite  is  a  carbonate  which  contains 
magnesium  and  calcium? 

4.  How  could  you  distinguish  powdered  chalk,  marble,  and  precipitated 
calcium  carbonate  from  one  another? 

5.  Write  equations  for  the  reactions  which  take  place  in  the  case  of  the  fol- 
lowing:  (a)  Ca3N2+H2O,    (6)  NaHCO3+  CaCl2,    (c)  CaCO3  +  C12  -f  H2O, 
(d)    NaOH  -j-  temporarily    hard    water,     (e)    NaOH  +  permanently    hard 
water,  and  (/)  NaOH + both  temporarily   and  permanently  hard  water,  (g) 
MgSO4  +  soap. 

6.  Write  equations  for  the  reactions  which  take  place    (a)  in  the  soften- 
ing of  water  by  permutit,  and   (6)  the  regeneration  of  the  latter. 

7.  (a)  What  weight  of  sodium  carbonate  is  required  to  soften  1,000,000 


550  INORGANIC  CHEMISTRY  FOR  COLLEGES 

gallons  of  permanently  hard  water  which  shows  20  degrees  of  hardness? 
(6)  If  a  water  is  temporarily  hard  (20  degrees)  what  weight  of  lime  is  re- 
quired to  soften  1,000,000  gallons  of  it? 

8.  (a)  Write  equations  for  the  reactions  by  which  tricalcium  phosphate 
can  be  converted  into  the  secondary  and  primary  salts,  and  free  phosphoric 
acid.     (6)  How  could  you  readily  distinguish  from  one  another  the  three 
salts? 

9.  (a)  Calculate  the  theoretical  percentage  of  available  chlorine  in  a  pure 
sample  of  bleaching  powder.     (6)  What  percentage  of  this  is  obtained  from 
commercial  bleaching  powder?     Devise  a  method  to  determine  (c)  the   total 
chlorine  and    (d)  the  available  chlorine  in  a  sample  of  bleaching  powder. 

10.  Calculate  the  molar  solubility  of  Ca(OH)2,  Sr(OH)2,  and  Ba(OH)2 
in  water  at  18°  and  state  the  normality  of  a  saturated  solution  of  each. 

11.  Starting  with  BaSO4  write  equations  for  the  reactions  used  in  making 
the  following:  (a)  BaCl?,   (6)  Ba(NO3)2,   (c)  Ba3(PO4)2,   (d)  BaO2. 


CHAPTER  XXXVII 
BERYLLIUM,    MAGNESIUM,   ZINC,   CADMIUM,   AND   MERCURY 

653.  The  elements  in  the  second  family  of  the  second  group  in 
the  periodic  classification  do  not  show  to  a  marked  degree  the 
gradation  in  physical  and  chemical  properties  which  is  ordinarily 
observed  in  the  case  of  other  families.     In  their  compounds  they 
are  bivalent,  but  mercury  forms  also  salts  in  which  the  metal  has 
the  valence  1.     The  hydroxides  of  the  metals  are  bases  which 
are  less  active  than  those  of  the  calcium  family,  and  the  salts 
derived  from  them  are  more  or  less  hydrolyzed  by  water.    The 
hydroxide  of  beryllium,  which  is  also  called  glucinum,  and  that  of 
zinc  are  weakly  acidic  and  dissolve  in  solutions  of  the  caustic 
alkalies. 

Beryllium  is  a  rare  element;  it  occurs  as  a  double  metasilicate 
with  aluminium  in  the  mineral  beryl,  Al2Be3(SiOs)6,  which  is 
called  emerald  when  colored  green  by  the  presence  of  a  small 
amount  of  chromium  silicate.  Beryllium  carbonate  dissolves  in 
water  and  is  highly  hydrolyzed  in  solution. 

MAGNESIUM 

654.  The  compounds  of  magnesium  are  widely  and  abun- 
dantly distributed  in  nature.     Dolomite,  a  double  carbonate  of 
magnesium  and  calcium,  occurs  as  a  rock  and  is  the  chief  constit- 
uent of  certain  mountain  ranges.     Magnesite  is  magnesium  car- 
bonate, talc  or  soapstone  (510)  is  a  hydrated  silicate  of  mag- 
nesium, and  asbestos  (510)  is  an  anhydrous  silicate.    The  sulphate 
and  chloride  of  magnesium  occur  in  sea  water  and  in  salt  beds. 

Compounds  of  magnesium  have  been  used  for  many  years. 
The  sulphate  occurs  in  the  water  of  the  mineral  springs  at  Epsom, 
England,  and  for  centuries  has  been  used  under  the  name  Epsom 
salt.  The  basic  carbonate,  "  magnesia  alba,"  was  known  to  the 

551 


552  INORGANIC  CHEMISTRY  FOR  COLLEGES 

alchemists.  The  free  element  was  first  isolated  in  an  impure 
condition  by  Davy,  and  was  prepared  in  1830  by  Liebig,  by  the 
action  of  potassium  on  fused  magnesium  chloride. 

Magnesium  is  manufactured  by  the  electrolysis  of  a  fused 
mixture  of  magnesium  and  potassium  chlorides,  which  are  obtained 
from  carnallite,  MgCl2,KCl,6H2O.  The  preparation  is  carried 
out  in  an  iron  vessel  which  serves  as  a  cathode,  the  anode  being 
a  rod  of  carbon. 

Magnesium  is  a  silver-white  metal,  which  is  characterized  by 
having  the  very  low  specific  gravity  1.75;  it  melts  at  650°,  boils 
at  1120°,  and  can  be  drawn  into  the  form  of  a  wire  and  rolled  into 
sheets.  The  metal  on  account  of  its  lightness  has  recently  been 
employed  in  making  alloys  to  be  used  in  the  construction  of  aero- 
planes (528).  Magnesium  is  less  active  than  the  metals  of  the 
calcium  family;  it  tarnishes  in  air  and  reacts  slowly  with  water. 
The  magnesium-mercury  couple  (545)  made  by  treating  the  metal 
with  a  little  mercuric  chloride  decomposes  water  fairly  rapidly. 

Magnesium  burns  in  the  air  with  an  intense  white  light  which 
is  rich  in  the  shorter  wave  lengths  of  the  visible  spectrum  (blue 
and  violet) ;  it  also  emits  the  so-called  ultra-violet  rays,  which  are 
still  shorter  in  wave  lengths,  are  not  visible,  but  affect  the  materials 
used  in  photographic  plates.  The  light  of  burning  magnesium  is 
said  to  be  about  sixty  times  as  active  chemically  as  that  produced 
by  burning  carbon.  In  photography  magnesium  is  burned  directly 
in  the  form  of  a  thin  sheet  or  as  a  powder.  In  the  latter  case  the 
finely  divided  metal  is  mixed  with  powdered  potassium  chlorate, 
the  amount  of  the  salt  used  being  that  required  to  furnish  just 
enough  oxygen  to  convert  the  metal  into  oxide.  When  such 
a  mixture  is  ignited  it  burns  instantaneously,  whereas  the  powder 
burns  sluggishly  and  soon  becomes  coated  with  oxide,  which 
stops  the  reaction.  When  magnesium  burns  in  the  air  a  part  of 
the  metal  is  converted  into  the  nitride,  MgsN2,  which  is  decom- 
posed by  boiling  water  with  the  formation  of  magnesium  hydroxide 
and  ammonia. 

655.  The  Oxide  and  Hydroxide  of  Magnesium. — Magnesium 
oxide  is  prepared  by  heating  magnesite,  MgCOs,  at  about  1700°; 
it  is  known  as  calcined  magnesia  and  is  obtained  as  a  white  porous 
powder,  which  is  infusible  and  does  not  react  with  water.  It  is 
employed  in  the  manufacture  of  fire-brick  that  are  used  in  lining 


BERYLLIUM,  MAGNESIUM,  ZINC,  CADMIUM,  AND  MERCURY  553 

furnaces  in  which  reactions  are  carried  out  at  very  high  tempera- 
tures. The  oxide  prepared  by  heating  the  precipitated  carbonate 
is  obtained  as  a  light  powder;  it  is  used  in  the  manufacture  of 
coverings  for  steam-pipes  as  an  insulator  to  prevent  the  transfer 
of  heat.  It  reacts  slowly  with  water  to  form  magnesium  hy- 
droxide. 

Magnesium  hydroxide,  Mg(OH)2,  is  made  by  treating  a  mag- 
nesium salt  with  a  solution  of  an  alkali.  It  is  very  slightly  soluble 
in  water,  about  0.01  gram  dissolving  in  1  liter,  but  the  concentra- 
tion of  a  saturated  solution  is  great  enough  to  permit  the  reaction  of 
the  base  with  ammonium  salts,  which,  accordingly,  dissolve  it: 

Mg(OH)2  +  2NH4C1  =  MgCl2  +  2NH3  +  2H20 

The  fact  that  the  hydroxide  is  not  precipitated  by  ammonium 
hydroxide  in  the  presence  of  ammonium  salts  has  already  been 
explained  at  some  length  (589) . 

656.  Magnesium  Chloride. — This  salt  resembles  calcium  chlo- 
ride in  composition  and  behavior.     It  crystallizes  from  water  as  a 
hexahydrate,  MgCl2,6H2O,  and  is  very  soluble  and  very  hygro- 
scopic.    Like  the  chloride  of  calcium  it  cannot  be  dehydrated 
without  decomposition,  hydrolysis  taking  place  even  to  a  greater 
extent  than  in  the  case  of  the  calcium  salt.     Anhydrous  mag- 
nesium chloride  can  be  prepared  by  first  dehydrating  a  double 
salt  of  the  composition  NH4Cl,MgCl2,6H2O,  and  then  driving  off 
the  ammonium  chloride  at  a  high  temperature.     The  hydrolysis  of 
magnesium  chloride  in  boiling  water  is  appreciable, 

MgCl2  +  2H20  ^  Mg(OH)2  +  2HC1 

and  for  this  reason  sea-water  and  other  water  which  contains  this 
salt  cannot  be  used  in  boilers. 

657.  Magnesium  Sulphate. — A  number  of  hydrates  of  this 
salt  are  known;   the  one  having  the  composition  MgSO4,7H2O  is 
obtained  when  a  solution  of  magnesium  sulphate  is  evaporated  so 
that   crystals   are  formed  from   it   at   room-temperature.     It  is 
efflorescent  and  very  soluble  in  water.     It  is  used  in  dyeing,  cotton 
printing,  and  other  industries,  and  in  medicine  as  a  purgative. 
The    monohydrate    of   magnesium    sulphate,    MgSO4,H2O,    ap- 
proaches in  solubility  that  of  gypsum,  CaS04,2H2O,  being  almost 


554  INORGANIC  CHEMISTRY  FOR  COLLEGES 

insoluble  in  water;   it  occurs  in  the  salt  beds  at  Stassfurt  as  the 
mineral  kieserite. 

658.  Magnesium  Carbonate. — This  compound  occurs  as  the 
mineral  magnesite,  MgCOs,  and  as  a  constituent  of  dolomite  (626). 
It  is  not  formed  in  the  pure  condition  by  precipitation  when  a 
soluble   carbonate  is  added  to   a  magnesium  salt,   the  product 
obtained  being  either  a  mixture  of  magnesium   carbonate  and 
magnesium  hydroxide  or  a  basic  salt  of  the  metal.     The  product 
obtained  by  the  precipitation  is  called  "  magnesia  alba,"  and  is 
used  in  medicine  and  as  a  cosmetic. 

659.  Boiler  Scale. — The  composition  of  hard  water  and  the 
methods  employed  in  its  treatment  when  it  is  to  be  used  for 
industrial  purposes  have  been  discussed  (629).     It  was    stated 
that  it  should  not  be  used  in  boilers  for  the  production  of  steam; 
this  is  due  to  the  fact  that  it  produces  on  the  boiler  tubes  a  deposit 
which  is  a  poor  conductor  of  heat,  and  the  fact  that  acids  are 
formed  which  corrode  the  metal  of  which  the  boiler  is  constructed. 

The  solids  which  separate  in  a  boiler  when  hard  water  is  evap- 
orated to  make  steam,  are  produced  either  as  a  porous,  non-adher- 
ing material,  which  can  be  easily  removed  by  "  blowing  off "  or 
washing  out  the  boiler,  or  as  a  dense  solid  which  coats  the  tubes 
and  sides  and  forms  what  is  called  "  boiler  scale."  In  the  first 
case  the  material  is  produced  largely  as  the  result  of  the  con- 
version of  the  soluble  bicarbonates  present  into  the  carbonates, 
which  precipitate.  The  scale  is  composed  principally  of  calcium 
sulphate  and  the  monohydrate  of  magnesium  sulphate,  MgSO^EbO 
(657) ,  which  is  practically  insoluble  and  is  formed  from  the  soluble 
sulphate  at  the  temperature  obtained  in  the  boiler.  Boiler  scale- 
also  contains  magnesium  hydroxide  which  is  present  in  it  as  the 
result  of  the  hydrolysis  of  the  magnesium  chloride  dissolved  in 
the  water.  It  will  be  recalled  that  calcium  sulphate  is  less 
soluble  in  hot  water  than  in  cold  water;  the  solubility  drops  off 
rapidly  as  the  water  is  heated  to  the  high  temperatures  obtained 
in  a  boiler.  At  68°  F.  water  dissolves  140.6  grains  of  the  salt  per 
gallon;  at  212°  F.  the  solubility  is  reduced  to  125.9  grains,  at  284° 
F.  when  the  steam  pressure  is  37  pounds  the  solubility  is  45.6 
grains,  and  at  356°  F.  which  corresponds  to  a  steam  pressure  of 
131  pounds  it  is  10.5  grains  per  gallon. 

The  heat  conductivity  of  the  compounds  present  in  boiler  scale 


BERYLLIUM,  MAGNESIUM,  ZINC,  CADMIUM,  AND  MERCURY  555 

is  small  compared  with  that  of  iron;  that  of  calcium  sulphate  has 
been  found  to  be  about  one-fiftieth  that  of  the  metal.  As  a  result, 
when  a  boiler  the  surface  of  which  is  covered  with  scale  is  used  to 
make  steam,  it  becomes  overheated  and  the  metal  may  get  red  hot 
and  become  deformed.  As  the  tubes  become  filled  with  scale  the 
surface  exposed  to  the  water  grows  smaller  and  the  steam-producing 
efficiency  of  the  boiler  rapidly  falls  off. 

The  corrosion  of  boilers  by  hard  water  is  due  chiefty  to  the 
presence  of  magnesium  chloride,  which  hydrolyzes  appreciably  at 
the  high  temperatures  reached.  The  acid  formed  dissolves  the 
iron  of  the  boiler. 

660.  Tests  for  Magnesium. — The  fact  that  sodium  carbonate 
precipitates  a  colorless,  insoluble,  basic  carbonate,  which  is  not 
formed  in  the  presence  of  ammonium  salts,  serves  to  separate 
magnesium  from  the  metals  of  the  alkaline  earths.     Magnesium- 
ammonium  phosphate,  MgNH4PO4,OH2O,  is  formed  as  a  crys- 
talline precipitate  when  disodium  phosphate,  ammonium  hydrox- 
ide, and  ammonium  chloride  are  added  to  a  solution  of  a  mag- 
nesium salt.     The  chloride  is  added  to  prevent  the  precipitation 
of  magnesium   hydroxide    (589)    by  the   ammonium  hydroxide. 
Magnesium-ammonium   phosphate   is   formed   and   used   in   the 
quantitative    determination    of    magnesium    or    of    phosphates. 
When  it  is  ignited  it  is  converted  into  magnesium  pyrophosphate, 
Mg2P20r.     Magnesium  salts  give  no  color  to  the  Bunsen  flame. 

ZINC 

661.  The  chief  ores  of  zinc  are  smithsonite  or  zinc  spar,  ZnCOs, 
and  sphalerite  or  zinc  blende,  ZnS;   zinc  silicate,  which  occurs  as 
calamine,    Zn2Si04,H2O,    and    franklinite,    Zn(FeO2)2,    are    also 
used.     The  metal  is  obtained,  after  roasting  the  ore,  by  reduction 
with  finely  divided  coal  in  cylindrical  fire-clay  retorts,  4  to  5  feet 
long  and  8  to  10  inches  in  diameter,  which  are  heated  by  gas. 
Since  reduction  of  the  oxide  takes  place  at  a  temperature  above  the 
boiling-point  of  the  metal,  the  vapor  of  the  latter  which  issues 
from  the  retort  is  led  through  a  condenser  made  of  fire-clay,  where 
it  is  deposited  in  the  molten  state.     At  first  a  mixture  of  zinc  oxide 
and  metallic  zinc  in  the  form  of  a  fine  powder,  which  is  known  as 
zinc  dust,  is  deposited ;  it  contains  about  10  per  cent  zinc  oxide  and 
on  account  of  the  fine  state  of  the  metal  finds  important  uses  as  a 


556  INORGANIC  CHEMISTRY  FOR  COLLEGES 

reducing  agent.  The  liquid  zinc  drawn  off  from  time  to  time  is 
cast  in  molds  and  is  called  "  spelter  ";  it  usually  contains  lead, 
cadmium,  and  iron,  and  smaller  quantities  of  arsenic,  antimony, 
and  sulphur.  It  is  refined  by  melting  in  a  furnace;  the  lead  pres- 
ent combines  with  a  little  zinc  and  settles  to  the  bottom,  and  on 
this  is  deposited  a  zinc-iron  alloy.  The  purest  zinc  is  made  by 
distillation  at  as  low  a  temperature  as  possible. 

Zinc  melts  at  419°,  boils  at  940°,  and  has  the  specific  gravity 
7.08  when  cast.  Cast  zinc  is  brittle,  but  when  heated  at  120° 
it  can  be  rolled  out  into  sheets,  which  remain  pliable  when  cooled. 
Just  above  200°  it  becomes  so  brittle  it  can  be  powdered.  The 
so-called  "  granulated  "  zinc  used  in  the  laboratory  is  prepared 
by  dropping  the  molten  metal  into  water. 

In  moist  air  zinc  becomes  covered  with  a  thin  coating  of  basic 
carbonate,  which  protects  it  from  further  action  (550).  It  burns 
in  air  when  heated,  reacts  with  most  non-metals,  forms  alloys 
with  other  metals  (542),  decomposes  steam  at  moderately  high 
temperatures,  and  reacts  with  hot  solutions  of  alkalies  (554). 
The  effect  of  impurities  in  the  metal  on  its  solution  in  acids  has 
been  described  (567). 

Zinc  is  used  for  roofs,  gutters,  etc.,  in  the  manufacture  of  brass 
and  other  alloys,  and  as  a  protective  coating  for  iron  which  is 
"galvanized"  or  "  sherardized "  (551).  Large  quantities  are 
consumed  in  dry  cells  (568).  The  metal  contracts  but  little  in 
passing  from  the  liquid  to  the  solid  state  and  can  be  cast;  so-called 
French  bronzes  are  cast  zinc,  which  is  electroplated  with  bronze 
or  brass. 

662.  Zinc  Oxide. — This  compound  is  used  extensively  under 
the  name  zinc  white  or  Chinese  white.  It  is  manufactured  by 
heating  metallic  zinc  and  conducting  the  vapor  of  the  metal 
through  a  flue  to  which  air  is  admitted;  the  oxide  formed  from  the 
burning  metal  is  collected  after  cooling  by  passing  it  along  with 
the  excess  of  air  and  nitrogen  through  settling  chambers  and  large 
bags  of  cotton  cloth.  At  times  the  vapor  of  the  metal  is  obtained 
directly  as  the  result  of  the  reduction  of  the  ore  with  coal. 

Zinc  oxide  turns  yellow  when  hot,  but  regains  its  whiteness  on 
cooling.  It  is  used  as  a  pigment  in  making  paint  and  as  a  filler 
in  the  manufacture  of  automobile  tires.  On  account  of  its  anti- 
septic properties  it  is  used  in  ointments. 


BERYLLIUM,  MAGNESIUM,  ZINC,  CADMIUM,  AND  MERCURY  557 

663.  Zinc  Hydroxide. — The  hydroxide  of  zinc,  which  is  formed 
on  the  addition  of  a  soluble  hydroxide  to  a  zinc  salt,  shows  both 
basic  and  acidic  properties.     Its  salts  are  hydrolyzed;    zinc  sul- 
phate shows  an  acidic  reaction  and  sodium  zincate,  Na2ZnO2,  a 
strong  basic  reaction.     In  order  to  convert  zinc  hydroxide  into  the 
latter  salt  an  excess  of  sodium  hydroxide  must  be  used  over  that 
indicated  by  the  following  reaction; 

Zn(OH)2  +  2NaOH  ^  Zn(ONa)2  +  2H2O 

Zinc  hydroxide  dissolves  in  a  solution  of  ammonia  as  the 
result  of  the  formation  of  the  compound  Zn(NH3)4(OH)2  (595). 

664.  Zinc  Chloride. — A  number  of  hydrates  of  zinc  chloride 
are  known,  all  of  which  are  deliquescent.     The  salt  is  usually 
sold  in  the  anhydrous   condition;   it  melts   at  290°  and  boils  at 
730°,  and  is  hydrolyzed  by  water.     When  mixed  with  zinc  oxide 
and  a  little  water  it  forms  a  plastic  mass  which  sets  to  a  solid: 

ZnCl2  +  ZnO  =  Zn2OCl2 

The  product  is  used  as  a  cement  and  for  other  purposes.     Zinc 
chloride  is  used  as  a  flux  in  soldering. 

665.  Other  Salts  of  Zinc. — Zinc  carbonate  is  formed  as  a  pre- 
cipitate when  sodium  bicarbonate  is  added  to  a  solution  of  a  zinc 
salt;    sodium  carbonate  which  is  slightly  hydrolyzed  produces  a 
basic  salt — Zn(OH)2,ZnCO3.    Zinc  sulphate,  ZnSO4,7H2O,  is  some- 
times called  white  vitriol;    it  is  isomorphous  with  green  vitriol, 
FeSO4,7H2O,  and  the  sulphates  of  other  bivalent  elements  which 
crystallize  with  7  molecules  of  water.     Zinc  sulphide  is  a  con- 
stituent of  lithophone  (649). 

666.  Test  for  Zinc  Salts. — The  sulphide  of  zinc,  ZnS,  is  formed 
as  a  white  precipitate  when  ammonium  sulphide  is  added  to  a 
solution  of  a  zinc  salt  or  a  zincate.     It  is  not  precipitated  by 
hydrogen  sulphide  in  the  presence  of  strong  acids,  in  which  it  is 
soluble.     Zinc  hydroxide  is  precipitated  by  sodium  hydroxide  and 
dissolves  in  an  excess  of  the  reagent;  it  dissolves  also  in  ammonium 
hydroxide.     In  determining  zinc  quantitatively  it  is  usually  pre- 
cipitated as  the  carbonate,  which  is  ignited  and  weighed  as  oxide. 


558  INORGANIC  CHEMISTRY  FOR  COLLEGES 


CADMIUM 

667.  Cadmium  is  obtained  as  a  by-product  in  zinc  smelting. 
The  small  amount  of  the  metal  present  in  the  ores  is  found  in  the 
zinc  dust  which  is  first  condensed  from  the  vapors  as  they  issue 
from  the  retorts  (661).     The  metal  is  more  volatile  than  zinc 
and  is  separated  by  distillation  from  the  zinc  dust  after  carbon 
has  been  added  to  reduce  the  oxides  present. 

Cadmium  is  a  silver-white  metal  which  melts  at  321°  and 
boils  at  778°.  It  dissolves  in  acids  less  readily  than  zinc.  It  is 
used  as  a  constituent  of  low-melting  alloys  (492) ,  and  in  amalgams 
for  dental  use.  Cadmium  sulphide,  which  is  bright  yellow,  is 
employed  as  an  artist's  color;  the  bromide  and  iodide  are  used  in 
precipitating  the  silver  halides  that  are  the  active  constituents  of 
photographic  plates  and  films. 

668.  Compounds    of    Cadmium. — A    large    number    of    com- 
pounds of  the  metal  are  known;   only  a  few  need  be  mentioned 
here.     The  chloride,  CdCl2,2H2O,  is  efflorescent  and  is  not  hydro- 
lyzed  by  water.     The  iodide,  Cdl2,  is  but  slightly  ionized  in  solu- 
tion— a  fact  thought  to  be  due  to  the  formation  of  a  salt  that  is 
composed  of  two  molecules  of  the  iodide,  to  which  the  formula 
Cd(Cdl4)  is  given.    This  view  is  a  reasonable  one  since  cadmium 
is  characterized  by  its  ability  to  form  stable  complex  salts.    Among 
these  are  double  chlorides  and  double  cyanides.     Cadmium  sul- 
phate, 3CdSO4,8H20,  is  not  isomorphous  with  the  sulphates  of 
zinc  and  magnesium.     The  precipitated  carbonate  of  the  metal  is 
a  normal  salt. 

669.  Tests  for  Cadmium  Salts. — Cadmium  hydroxide  is  formed 
as  a  white  precipitate  when  a  solution  of  an  alkali  is  added  to  a 
cadmium  salt;  it  is  insoluble  in  alkalies,  but  dissolves  in  ammonium 
hydroxide  to  form  the  compound  Cd(NHs)4(OH)2.      The  sul- 
phide, which  is  yellow,  is  insoluble  in  dilute  solutions  of  strong 
acids,  and  is,  consequently,  precipitated  along  with  copper  and  the 
other  metals  of  the  so-called  second  group  in  qualitative  analysis. 
It  dissolves  in  strong  solutions  of  acids.     Cadmium  can  be  sep- 
arated from  copper  by  adding  potassium  cyanide  to  the  solution 
of  the  compounds  of  the  two  metals  and  passing  in  hydrogen  sul- 
phide.    The  soluble  complex  cyanide  of  copper  which  is  formed, 
KCu(CN)2,    is    not    decomposed,    whereas    the    cadmium    salt, 


BERYLLIUM,  MAGNESIUM,  ZINC,  CADMIUM,  AND  MERCURY  559 

K2Cd(CN)4,  is  converted  into  cadmium  sulphide,  which  appears 
as  a  yellow  precipitate. 

MERCURY 

670.  On  account  of  its  occurrence  in  the  free  condition  mer- 
cury has  been  known  for  over  2000  years.     The  fact  that  it  is  a 
liquid  and  resembles  silver  in  its  properties  led  to  the  Latinized 
name  of  the  element,  hydrargyrum,  which  was  derived  from  the 
Greek  words  signifying  liquid  silver.     Mercury  played  an  impor- 
tant part  in  the  early  theories  of  the  constitution  of  matter.  It  was 
named  by  the  alchemists  after  Mercury,  the  messenger  of  Jove, 
on  account  of  its  rapid  disappearance  when  heated,  and  was  indi- 
cated in  alchemical  books  by  the  symbol  which  represented  a 
messenger's  wand — the  same  symbol  as  that  used  in  astrology  to 
represent   the   planet  Mercury.     The   salts   of  mercury  have   a 
marked  effect  on  the  human  body,  and  were  studied  in  consider- 
able detail  in  the  sixteenth  century  by  the  iatrochemists,  who 
investigated  the  action  of  inorganic  substances  as  drugs. 

Mercury  forms  two  classes  of  compounds;  the  members  of  one 
of  these,  in  which  the  metal  is  bivalent,  resemble  to  some  extent 
the  analogous  compounds  of  cadmium  and  zinc;  those  of  the 
other  class  contain  mercury  with  the  valence  1.  The  compounds 
of  univalent  mercury  possess  properties  analogous  to  those  of  the 
compounds  of  other  heavy  metals,  copper  and  silver,  when  the 
latter  have  the  valence  1 ;  they  are  called  mercurous  compounds, 
and  those  derived  from  bivalent  mercury  are  mercuric  com- 
pounds. 

671.  Preparation  and  Properties  of  Mercury. — The  element 
is  found  in  the  free  condition  but  is  obtained  chiefly  from  cinnabar, 
HgS,  which  occurs  as  a  red  crystalline  mineral  in  California,  Spain, 
Italy,  and  Austria.     The  ground  mineral  is  heated  in  a  stream 
of  hot  air  which  oxidizes  it — HgS  +  62  =  Hg  +  S(>2 — and  the 
vapor  of  the  metal  formed  is  condensed.     It  can  be  purified  by 
distillation,  or  by  bringing  it  into  contact  with  dilute  nitric  acid, 
which  dissolves  the  impurities  present.     The  latter  process  is 
carried  out  in  the  laboratory  by  allowing  the  metal  to  fall  in 
minute  globules  through  the  acid  contained  in  a  long  glass  tube. 
It  can  also  be  purified  by  bubbling  air  through  it  for  a  long  time. 


560  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  metals  dissolved  in  the  mercury  are  oxidized  and  form  on  its 
surface  a  scum,  which  can  be  removed  by  filtering  the  liquid 
through  chamois  skin. 

Mercury  is  the  only  metal  which  is  liquid  at  ordinary  tempera- 
tures; it  melts  at  -38.9°,  boils  at  357°,  and  has  the  specific  grav- 
ity 13.6  at  0°.  It  is  used  in  barometers,  thermometers,  and  air 
pumps  for  laboratory  use,  as  a  solvent  for  gold  and  silver  in  extract- 
ing these  metals  from  their  ores,  as  the  cathode  material  in  the 
electrolytic  preparation  of  the  caustic  alkalies,  in  amalgamating 
zinc  to  be  used  in  electric  batteries  (567),  and  in  the  preparation 
of  amalgams  with  tin,  silver  or  gold  which  are  used  by  dentists  to 
fill  cavities  in  teeth.  Finely  divided  mercury  which  has  the 
appearance  of  a  powder  is  used  in  medicine  in  the  form  of  pills. 

Mercury  unites  with  oxygen  when  it  is  heated  with  it  at  about 
300°  and  forms  mercuric  oxide.  It  combines  with  the  halogens  and 
sulphur,  does  not  react  with  dilute  solutions  of  non-oxidizing  acids, 
but  dissolves  in  nitric  acid.  It  decomposes  hydrogen  sulphide 
and  is  converted  into  black  mercuric  sulphide.  Mercury  dissolves 
many  metals,  with  some  of  which  it  unites  chemically.  The 
compound  of  mercury  and  sodium  having  the  composition  NaHg2 
is  a  solid. 

672.  The   Oxides  of  Mercury. — Mercurous  oxide,   Hg2O,   is 
formed  as  a  black  precipitate  when  sodium  hydroxide  is  added  to  a 
solution  of  a  mercurous  salt;  it  is  slowly  decomposed  into  mercury 
and  oxygen  by  light  or  by  gentle  heating.     Mercuric  oxide,  HgO, 
is  formed  in  a  similar  way  from  mercuric  salts;  under  these  con- 
ditions it  appears  as  a  yellow,  amorphous  powder  which  turns 
dark  brown  when  heated.     When  the  oxide  is  prepared  by  heating 
mercury  at  300°-350°,  or  by  calcining  mercuric  nitrate,  Hg(NO3)2, 
it  is  a  bright  red  powder.     No  hydroxides  of  mercury  exist;  the 
reactions  which  would  lead  to  their  formation  produce  the  oxides 
of  the  metal. 

673.  The  Nitrates  of  Mercury. — If  mercury  is  dissolved  in  hot 
concentrated  nitric  acid  it  is  converted  into  mercuric  nitrate, 
Hg(NO3)2,8H2O;  when  left  in  contact  with  the  dilute  acid  it  reacts 
slowly  to  form  mercurous  nitrate,  which  crystallizes  from  water  as  a 
monohydrate,  HgNO3,H2O.     Both  salts  are  hydrolyzed  by  water 
and  form  insoluble  basic  salts;  they  remain  in  solution  only  in  the 
presence  of  an  excess  of  nitric  acid.     Mercuric  nitrate  can  be  con- 


BERYLLIUM,  MAGNESIUM,  ZINC,  CADMIUM,  AND  MERCURY  561 

verted  into  mercurous  nitrate  by  shaking  its  solution  with  mer- 
cury: 

Hg(N03)2  +  Hg  =  2HgN03 

674.  The   Halides   of   Mercury. — Mercuric   chloride,    HgCl2, 
which  is  known  as  corrosive  sublimate,  is  made  by  subliming  a 
mixture  of  mercuric  sulphate  and  sodium  chloride: 

HgSO4  +  2NaCl  =  HgCl2  +  Na2S04 

It  boils  at  307°,  and  when  the  mixture  is  heated  to  this  temperature 
the  conditions  are  such  that  the  double  decomposition  takes  place. 
Mercuric  chloride  unites  directly  with  protein  material  and  is,  as  a 
consequence,  a  poison.  When  taken  into  the  body  it  interferes 
with  the  proper  functioning  of  the  kidneys,  and  the  waste  products 
formed  as  the  result  of  the  normal  changes  which  take  place,  are 
not  excreted  and  may  finally  produce  death.  The  antidote  com- 
monly used  for  corrosive  sublimate  is  white  of  egg,  which  is  a  pro- 
tein. If  this  is  administered  soon  enough  it  unites  with  the  salt 
to  form  an  insoluble  compound,  which  can  be  removed  from  the 
stomach  by  means  of  an  emetic.  A  very  dilute  solution  of  mer- 
curic chloride  is  used  as  an  antiseptic  in  surgery. 

Mercurous  chloride,  HgCl,  is  used  in  medicine  under  the  name 
calomel.  It  is  prepared  by  subliming  a  mixture  of  mercuric  chlo- 
ride and  mercury — HgCl2  +  Hg  =  2HgCl — or  a  mixture  of 
mercuric  sulphate,  mercury,  and  salt.  It  is  used  in  medicine  as  a 
purgative.  Mercurous  chloride  is  formed  as  a  white  precipitate 
wrhen  a  soluble  chloride  is  added  to  a  solution  of  a  mercurous  salt. 
It  is  also  readily  formed  by  reducing  mercuric  chloride;  when 
stannous  chloride,  SnCl2,  is  used,  it  is  oxidized  to  stannic  chloride, 
and  mercurous  chloride  or  mercury  is  formed  depending  upon  the 
quantity  of  the  reducing  agent  used : 

2HgCl2  +  SnCl2  =  2HgCl  +  SnCU 
2HgCl  +  SnCl2  =  2Hg  +  SnCU 

The  mercury  appears  as  a  black  precipitate. 

675.  Mercuric  iodide,  HgI2,  is  obtained  as  a  scarlet  precipitate 
when  a  soluble  iodide  is  added  to  a  solution  of  a  mercury  salt. 
When  the  salt  is  heated  above  128°  it  is  converted  into  a  form  which 
is  yellow,  melts  at  223°,  and  on  cooling  yields  yellow  crystals. 


562  INORGANIC  CHEMISTRY  FOR  COLLEGES 

When  these  are  cooled  to  room  temperature  they  are  unstable  and 
change  on  being  touched  with  a  sharp  object  into  the  scarlet 
variety.  Mercuric  iodide  dissolves  in  an  excess  of  potassium 
iodide  to  form  a  stable,  colorless,  complex  salt  of  the  composition 
K2HgI4,  from  which  the  mercury  is  not  precipitated  by  bases. 

Mercurous  iodide,  Hgl,  is  formed  as  a  dark-green  solid  when 
iodine  is  rubbed  with  an  excess  of  mercury;  when  a  soluble 
iodide  is  added  to  a  solution  of  a  mercurous  salt  it  is  precipitated 
as  a  greenish-yellow  powder,  which  decomposes  spontaneously 
into  mercuric  iodide  and  mercury:  2HgI  =  Hg  +  Hgl2. 

676.  When  a  solution  of  ammonia  is  added  to  mercury  salts 
the  reactions  which  take  place  are  of  a  different  nature  from  any 
that  have  been  described;    the  so-called  ammono-compounds  are 
formed,  of  which  those  derived  from  the  halides  are  of  the  most 
importance  because  they  are  used  in  analytical  chemistry. 

The  addition  of  a  solution  of  ammonia  to  one  of  mercuric 
chloride  causes  the  precipitation  of  a  white  compound  of  the 
formula  Hg(NH2)Cl  which  is  called,  "infusible  white  precipitate" 
or,  better,  ammono-basic  mercuric  chloride: 

HgCl2  +  2NH3  =  Hg(NH2)Cl  +  NH4C1 

The  interaction  of  ammonia  to  form  a  compound  of  this  type  as 
the  result  of  the  replacement  of  a  part  of  the  acid  radical  by  the 
NH2  group  resembles  closely  reactions  of  hydrolysis;  in  the  latter 
case  water  separates  into  H  and  OH,  the  hydrogen  unites  with  the 
acid  radical,  and  the  OH  takes  the  place  of  the  latter  and  a  basic 
salt  is  formed : 

HgCl2  +  HOH  =  Hg(OH)Cl  +  HC1 

in  the  case  of  ammonia  the  separation  of  the  latter  into  H  and  NH2 
takes  place  and  the  hydrogen,  as  before,  unites  with  a  part  of  the 
acid  radical,  which  is  replaced  by  the  NH2  group.  On  account  of 
this  close  analogy  the  reaction  is  said  to  be  one  of  ammonolysis 
and  the  salt  formed  is  called  an  ammono-basic  salt. 

677.  When  mercurous  chloride  is  treated  with  a  solution  of 
ammonia,  it  decomposes  into  mercury  and  mercuric  chloride  and 
the  latter  reacts  with  the  ammonia  to  form  ammono-basic  mercuric 
chloride.     The  mercury  is  deposited  throughout  the  precipitate 
as  a  fine  powder  and,  as  a  result,  the  precipitate  is  black.    It  was 


BERYLLIUM,  MAGNESIUM,  ZINC,  CADMIUM,  AND  MERCURY  563 

this  reaction  that  gave  its  name  to  calomel,  the  word  being  derived 
from  the  Greek  words  which  signify  "  beautiful  black." 

678.  In  the  case  of  mercuric  iodide  the  reaction  is  more  com- 
plex; it  is  probable  that  the  compound  Hg(NH2)I  is  first  formed 
and  then  reacts  with  a  second  molecule  of  mercuric  iodide  as 
follows: 

Hg(NH2)I  +  HgI2  -  Hg(NHg)I  +  2HI 

The  two  hydrogen  atoms  of  the  NH2  group  unite  with  the  two 
iodine  atoms  in  mercuric  iodide  to  form  hydrogen  iodide,  and  are 
replaced  by  a  mercury  atom;  the  graphic  formula  of  the  compound 
is,  accordingly,  Hg  =  N  — Hg  — I.  The  reaction  is  brought  about 
by  adding  ammonia  to  a  solution  of  the  iodide  in  potassium  iodide 
(K2Hgl4).  The  substance  is  a  brown  precipitate  and  has  high 
tinctorial  power,  that  is,  an  exceedingly  small  amount  of  it  can 
be  recognized  by  its  color. 

A  solution  of  mercuric  iodide  in  potassium  iodide  to  which 
potassium  hydroxide  has  been  added  is  known  as  Nessler's  reagent 
and  is  used  to  detect  small  amounts  of  ammonia  in  water  analysis. 
If  the  water  contains  ammonium  salts  the  base  present  in  the 
solution  liberates  ammonia,  which  then  interacts  with  the  mer- 
curic iodide  to  produce  a  brown  color.  The  amount  of  ammonia 
present  in  the  solution  is  determined  by  matching  the  shade  pro- 
duced by  it  with  that  formed  in  one  of  a  series  of  standards  pre- 
pared from  the  reagent  and  solutions  containing  known  amounts 
of  ammonia.  This  is  an  example  of  a  so-called  "  colorimetric  " 
method  of  analysis,  many  of  which  are  used  in  quantitative  analysis. 

679.  Sulphides    of    Mercury. — When    hydrogen    sulphide    is 
passed  into  a  solution  of  a  mercurous  salt  at  -ordinary  tempera- 
tures, a  mixture  of  mercury  and  mercuric  sulphide,    HgS,    is 
obtained.     Mercurous  sulphide,  Hg2S,  is  stable  only  at  low  tem- 
peratures.    Mercuric  sulphide  is  formed  as  a  black  precipitate 
when  hydrogen  sulphide  is  passed  into  a  mercuric  salt.     It  is 
very  insoluble  in  water  and  acids  and  is  not  attacked  by  boiling 
nitric  acid,  its  inertness  in  this  respect  being  utilized  in  separating 
it  from  the  sulphides  of  other  metals  precipitated  in  the  course  of  a 
qualitative  analysis.     It  is  soluble  in  aqua  regia. 

The  precipitated,  black  sulphide  is  converted  into  the  bright 
red  crystalline  variety  when  it  is  sublimed.  The  sulphide,  which 


564  INORGANIC  CHEMISTRY  FOR  COLLEGES 

occurs  as  cinnabar,  was  formerly  much  used  as  a  red  pigment  under 
the  name  vermilion.  The  material  used  at  present  is  manufac- 
tured either  by  heating  a  mixture  of  mercury  and  sulphur  until 
the  black  sulphide  is  formed,  and  then  subliming  the  product,  or 
by  heating  a  mixture  of  mercury  and  sulphur  with  a  strong  solu- 
tion of  caustic  potash. 

680.  Other  Salts  of  Mercury. — Mercuric  cyanide,  Hg(CN)2,  is 
prepared  from  precipitated  mercuric  oxide  and  hydrocyanic  acid, 
HCN;  it  forms  prismatic  crystals  that  decompose  on  heating  into 
mercury  and  cyanogen,  (CN)2,  a  poisonous  gas,  which  burns  with  a 
characteristic  flame  composed  of  a  pink  cone  surrounded  by  a  blue 
mantle.     Mercuric  fulminate,  Hg(ONC)2,  is  formed  when  mercury 
is  dissolved  in  nitric  acid  and  alcohol  is  added  to  the  solution.     It 
explodes  rapidly  when  struck  and  is  used  in  detonating  caps  to 
explode  gunpowder,  nitroglycerine,  etc. 

681.  Tests  for  Mercury  Salts. — Mercurous  salts  are  precip- 
itated  by   hydrochloric   acid,    and   when   the   white   mercurous 
chloride  formed  is  treated  with  ammonia  it  turns  black  (676). 
Mercuric   salts   are   precipitated   by   hydrogen   sulphide   in   the 
presence  of  hydrochloric  acid.     The  black  sulphide  is  separated 
from  other  sulphides  by  treatment  with  hot  nitric  acid,  in  which 
it  alone  does  not  dissolve.     It  is  then  dissolved  in  aqua  regia,  and 
the  solution  of  the  chloride  is  treated  with  stannous  chloride,  which 
precipitates  white  mercurous  chloride  and  finally  mercury   as  a 
black  powder  (674).     All  the  salts  of  mercury  are  volatile  when 
heated. 

EXERCISES 

1.  How  could  you  prepare  MgSO4,H2O  from  MgSO4,  7H2O? 

2.  Explain  why  Na2CO3  precipitates  CaCO3  from  calcium  salts  and  a 
basic  carbonate  from  magnesium  salts. 

3.  A  sample  of  a  mineral  containing  only  MgCOs  and  CaCOs,  weighing 
1  gram,  was  analyzed  for  Mg.     The  Mg2P2O7  obtained  weighed  0.6032  gram. 
Calculate    (a)  the  weight  of  MgCO3  in  the  sample,    (6)  the  percentage  of 
Mg  in  the  mineral,  (c)  the  weight  of  CO2  that  could  be  obtained  from  1  gram 
of  the  mineral. 

4.  Would  the  addition  of  calcium  hydroxide  to  a  water  which  caused  the 
corrosion  of  a  boiler  be  advantageous?     Why? 

5.  Show  by  ionic  equations  why  the  neutral  carbonate  of  zinc  is  pre- 
cipitated by  NaHCOs  and  the  basic  carbonate  by 


BERYLLIUM,  MAGNESIUM,  ZINC,  CADMIUM,  AND  MERCURY  565 

6.  (a)  Write  ionic  equations  to  represent  the  hydrolysis  of  MgS  and  CaS. 
(b)  In  which  case  is  the  hydrolysis  greater?     Give  a  reason  for  your  answer. 

7.  Write  equations  for  reactions  by  means  of  which    (a)  magnesium  and 
zinc  salts  can  be  distinguished  in  three  different  ways,  and    (6)  zinc  salts 
from  cadmium  salts  in  two  ways. 

8.  What  chemical  reactions  would  occur  if  a  sample  of  mercury  contain- 
ing small  amounts  of  zinc  and  lead  were  agitated  with  a  solution  of  mercurous 
nitrate?     For  what  purpose  could  the  process  be  used? 

9.  Write  equations  for  the  reactions  which  take  place  between  Hg  and 
(a)  concentrated  HNO3  and   (&)  dilute  HNO3. 

10.  Write  equations  for  reactions  by  which  Hg(NO3)2  can  be  converted 
into  (a)  HgCl,   (&)  HgO,   (c)  HgCl2,   (d)  HgNO3,   (e)  HgSO4. 


CHAPTER  XXXVIII 
ALUMINIUM 

682.  Aluminium,  which  is  sometimes  called  aluminum,  is 
the  typical  metal  in  the  third  group  in  the  periodic  classification 
of  the  elements.  With  the  exception  of  boron,  which  shows  acid- 
forming  properties  only,  the  other  elements  in  the  group  resemble 
aluminium  closely.  Thallium,  in  addition  to  being  trivalent, 
forms  a  number  of  compounds  in  which  it  shows  the  valence  1; 
these  resemble  closely  the  analogous  derivatives  of  silver. 

Other  metals  which  can  show  a  higher  valence  than  3  and, 
therefore,  fall  into  other  groups  in  the  periodic  classification,  form 
important  compounds  in  which  the  element  is  trivalent;  among 
these  are  chromium,  iron,  and  manganese.  The  compounds 
derived  from  these  metals  in  the  trivalent  condition  resemble 
closely  in  their  chemical  properties  the  analogous  compounds  of 
aluminium.  We  shall  learn,  therefore,  through  the  study  of 
aluminium  compounds  the  chemical  behavior  of  the  derivatives  of 
a  trivalent  metal. 

A  marked  difference  exists  between  the  chemical  properties  of 
the  compounds  of  the  univalent  metals  and  those  of  the  analogous 
compounds  of  the  bivalent  metals.  We  have  seen,  for  example, 
that  the  hydroxides  of  the  alkali  metals  are  strong  bases,  which 
are  very  soluble  in  water,  and  form  salts  that  are  not  hydrolyzed. 
The  carbonates  of  these  metals  resist  high  temperatures,  and  the 
salts  of  other  oxygen  acids  are  relatively  stable  when  heated. 
The  hydroxides  of  the  bivalent  metals  are  weaker  bases,  much  less 
soluble  in  water,  and  their  salts  are  hydrolyzed  to  a  small  extent. 
The  carbonates  of  these  metals  are  converted  by  heat  into  oxides, 
and  the  salts  of  other  oxygen  acids  are  much  more  readily  decom- 
posed than  those  of  the  alkali  metals. 

The  decreased  activity  of  the  metals  as  base-forming  elements 
when  they  are  in  the  trivalent  condition  is  marked.  The  hydrox- 

566 


ALUMINIUM  567 

ides  of  these  metals  are  very  weakly  base-forming,  and  the  salts 
derived  from  them  are  highly  hydrolyzed;  carbonates  cannot  be 
prepared  in  aqueous  solution  by  precipitation,  because  they  are 
completely  hydrolyzed  by  water.  The  salts  are  decomposed  by 
heat  at  comparatively  low  temperatures.  The  hydroxides  of 
some  of  the  metals  dissolve  in  solutions  of  strong  bases,  and  act, 
therefore,  as  acids.  A  comparison  of  the  behavior  of  the  deriva- 
tives of  the  three  types  of  metals  brings  out  clearly  the  fact 
already  mentioned,  that,  in  general,  low  valence  toward  negative 
atoms  or  groups  is  associated  with  base-forming  properties  and 
high  valence  with  acid-forming  properties. 

683.  The  occurrence  of  aluminium  has  already  been  noted.     It 
is  the  most  abundant  of  the  metals  and  the  most  widely  distributed. 
It  is  found  in  feldspars,  micas,  kaolin,  clay,  bauxite,  cryolite,  alun- 
ite,  corundum,  and  certain  gems  (510). 

Compounds  of  aluminium  have  been  known  for  many  years, 
and  they  were  recognized  as  being  derived  from  a  metal  that  had 
not  been  isolated.  Many  unsuccessful  attempts  were  made  to 
separate  it  from  its  oxide,  but,  finally,  Wohler  in  1827  obtained 
the  metal  as  the  result  of  the  action  of  potassium  on  fused  alu- 
minium chloride.  This  method  was  too  expensive  to  be  used  for 
commercial  purposes,  and  aluminium,  which  has  found  a  very 
important  place  in  industry,  was  not  manufactured  until  the 
electrolytic  method  devised  in  1886  by  Hall,  was  put  into  success- 
ful operation. 

684.  Preparation  and  Uses  of  Aluminium. — In  the  Hall  process 
the  metal  is  obtained  by  the  electrolysis  of  aluminium  oxide  dis- 
solved in  a  bath  of  fused  cryolite,  AlFs,3NaF.     The  material  is 
contained  in  rectangular  iron  pots  that  are  lined  with  thick  plates 
of  carbon,  which  serve  as  the  cathode.     The  anodes  are  large 
graphite  rods.     In  starting  the  reduction,  the  anodes  are  brought 
into  contact  with  the  bottom  of  the  tank  and  heat  is  developed 
by  the  current  as  the  result  of  the  poor  contact  established.     Cry- 
olite is  placed  in  the  cell,  and  when  the  bath  is  in  the  molten  con- 
dition the  anodes  are  withdrawn  from  the  cathode,  and  aluminium 
oxide  is  added.     The  resistance  of  the  cell  produces  enough  heat  to 
keep  the  material  in  the  fused  condition.     From  time  to  time 
molten  aluminium  is  withdrawn  from  the  bottom  of  the  cell,  and 
more  oxide  is  added;    the  process  is  a  continuous  one  because 


568  INORGANIC  CHEMISTRY  FOR  COLLEGES 

the  cryolite  is  not  decomposed  by  the  current  and  serves  only  as  a 
solvent  for  the  oxide. 

Aluminium  has  a  very  low  density,  2.6;  it  is  used  in  construc- 
tion when  a  metal  is  required  and  weight  is  an  important  consider- 
ation. It  is  ductile,  malleable,  and  can  be  rolled.  Its  tensile 
strength  is  low  in  comparison  with  that  of  iron;  it  cannot  be 
machined  and  polished  readily,  and  does  not  yield  good  castings. 
These  defects  can  be  overcome  by  alloying  it  with  other 
metals. 

Alloys  of  copper  and  aluminium  which  contain  from  5  to  10 
per  cent  of  the  latter  are  called  aluminium  bronzes.  They 
have  a  fine  yellow  color  resembling  gold  and  are  used  in 
making  imitation  jewelry  and  statuary.  The  alloys  which  con- 
tain from  11  to  89  per  cent  aluminium  are  brittle  and  highly 
crystalline  and  are  not  used.  Those  that  contain  90-93  per  cent 
aluminium  are  used  in  making  castings  and  are  silver-white. 
On  account  of  its  low  electrical  resistance  aluminium  is  used  in 
certain  cases  in  wires  and  cables  as  conductors.  In  a  finely  divided 
condition  it  is  used  in  making  a  paint,  and,  in  the  form  of  leaf,  for 
stamping  letters  on  book  covers. 

At  high  temperatures  aluminium  is  an  active  deoxidizing  agent. 
When  a  mixture  of  the  powdered  metal  and  a  finely  divided  oxide 
of  iron  is  ignited,  a  vigorous  reaction  takes  place  and  aluminium 
oxide  and  iron  in  the  molten  condition  are  formed.  The  tem- 
perature reached  is  about  2300°.  The  mixture,  which  is  called 
thermite,  is  used  in  welding  together  the  ends  of  steel  rails,  for 
mending  broken  shafts  of  marine  engines,  etc.  It  was  used  in 
incendiary  bombs  during  the  recent  war. 

Oxides  of  metals  which  do  not  react  with  carbon  at  the  temper- 
ature obtained  by  burning  coal,  are  reduced  when  heated  with  alu- 
minium. The  method  has  been  used  to  prepare  chromium  and 
similar  metals.  Large  quantities  of  aluminium  are  used  in 
deoxidizing  iron  and  steel;  the  metal  unites  with  the  gases  in  the 
metal,  and  castings  free  from  blow  holes  are  obtained. 

The  chief  physical  and  chemical  properties  of  the  metal  have 
already  been  described.  It  dissolves  in  hydrochloric  acid  but 
not  in  dilute  sulphuric  or  nitric  acid.  It  is  soluble  in  solutions  of 
the  caustic  alkalies  as  the  result  of  the  formation  of  aluniin- 
ates  (654). 


ALUMINIUM  569 

685.  Aluminium  Hydroxide. — When  a  salt  of  aluminium    is 
treated  with  a  solution  of  a  base,  aluminium  hydroxide  is  formed 
as  a  gelatinous  precipitate: 

A1CU  +  SNttiOH  =  A1(OH)3  +  3NEUC1 

The  hydroxide  dissolves  in  solutions  of  caustic  alkalies  as  the 
result  of  the  fact  that  it  shows  weakly  acidic  properties.  The 
salts  are  derived  from  a  dehydrated  form  of  the  hydroxide — 
A1(OH)3  -  H2O  =  HAlO-j— and  are  called  metaluminates: 

NaOH  +  A1(OH)3  =  NaAlO2  +  2H2O 

The  salt  is  highly  hydroly zed  and  remains  in  solution  only  in  the 
presence  of  an  excess  of  sodium  hydroxide.  When  an  acid  is 
added  cautiously  to  the  solution,  aluminium  hydroxide  is  repre- 
cipitated.  The  hydroxide  does  not  dissolve  in  ammonium  hydrox- 
ide, which  is,  accordingly,  used  when  it  is  desired  to  precipitate  it 
quantitatively. 

686.  Aluminium  Oxide. — Corundum,  AfeOs,  is  a  pure  crys- 
talline form  of  the  oxide  of  aluminium  which  occurs  as  a  mineral; 
it  stands  second,  next  to  diamond,  in  the  scale  of  hardness  (529). 
It  was  formerly  much  used  under  the  name  of  emery  for  polishing 
hard  surfaces,  but  has  been  largely  replaced  by  carborundum,  SiC, 
an  electric  furnace  product  (218).    Ruby  and  sapphire  are  crys- 
talline forms  of  the  oxide,  the  color  being  produced  probably  in  the 
case  of  the  former  by  a  trace  of  chromium  compounds  and  of  the 
latter  by  aluminates  of  iron  and  titanium.    Synthetic  rubies  and 
sapphires  are  now  manufactured  by  fusing  in  the  oxyhydrogen 
flame  aluminium  oxide  to  which  has  been  added  the  materials 
required  to  produce  the  proper  color.     Aluminium  oxide  is  often 
called  alumina. 

Aluminium  oxide  is  used  in  making  chemical  apparatus 
designed  to  resist  high  temperatures.  The  object  is  fashioned 
out  of  the  powdered  oxide  and  heated  in  an  electric  furnace  until 
the  outside  has  just  fused.  Although  the  oxide  of  aluminium  pre- 
pared by  dehydrating  the  hydroxide  is  readily  soluble  in  acids,  it 
becomes  very  inactive  after  fusion.  Ware  prepared  in  this  way  is 
sold  under  the  trade  name  alundum. 


570          'INORGANIC  CHEMISTRY  FOR  COLLEGES 

687.  Aluminium     Chloride. — When   a   solution   prepared    by 
dissolving  aluminium  or  its  oxide  or  hydroxide  in  hydrochloric 
acid  is  evaporated,  the  chloride  separates  in  crystals,  which  have 
the  formula  AlCl3,6H2O.     The  compound  is  completely  hydro- 
lyzed  when  an  attempt  is  made  to  dehydrate  it.     The  anhydrous 
salt,  which  is  much  used  in  the  preparation  of  certain  organic 
compounds,  is  prepared  by  the  action  of  chlorine  on  aluminium. 
It  has  a  vapor  pressure  of  760  mm.  at  183°,  and  sublimes  without 
melting,  giving  a  white  crystalline  solid. 

The  other  halides  of  aluminium  are  made  in  a  similar  way; 
they  all  form  with  the  halides  of  the  alkali  metals,  characteristic 
double  salts  of  which  cryolite,  AlFs,3NaF,  is  a  noteworthy 
example. 

688.  Aluminium    Sulphate.— This    salt,    A12(SO4)3,18H2O,    is 
extensively  used  in  the  industries  and  is  manufactured  in  large 
quantities  by  the  action  of  sulphuric  acid  on   clay,   bauxite,   or 
aluminium  hydroxide  prepared  from  cryolite. 

Aluminium  sulphate  is  used  in  purifying  water.  It  is  con- 
verted, after  being  dissolved,  into  the  hydroxide  by  the  addition 
of  slaked  lime,  if  the  water  is  soft,  or  by  the  calcium  bicarbonate 
present,  if  the  water  is  temporarily  hard.  The  hydroxide  is  formed 
as  a  gelatinous  precipitate  that  causes  the  finely  divided  suspended 
matter  to  coagulate  into  large  particles  which  settle  and  thus 
leave  the  water  clear.  The  precipitate  absorbs  most  of  the  bac- 
teria present. 

Aluminium  salts  are  used  in  mordanting  cotton  and  linen 
before  they  are  dyed.  These  fibers  do  not  absorb  from  solutions 
many  kinds  of  dyes  in  such  a  form  that  they  are  permanently 
held  and  not  removed  by  washing.  To  overcome  this,  the 
material  to  be  dyed  is  first  mordanted,  that  is,  an  insoluble 
substance  is  deposited  on  the  fiber,  which  unites  with  the  dye 
and  renders  it  insoluble  in  water.  When  aluminium  salts  are 
used,  the  fiber  absorbs  the  aluminium  hydroxide  formed  as  the 
result  of  their  hydrolysis,  and  the  hydroxide  unites  with  the  color- 
ing matter  of  the  d}7e-bath.  Wool  absorbs  readily  aluminium 
hydroxide  from  a  solution  of  the  sulphate,  but  since  cotton  and 
linen  do  not,  they  are  mordanted  in  a  bath  containing  a  soluble 
basic  sulphate  formed  by  dissolving  aluminium  hydroxide  in  a 
solution  of  the  sulphate.  Aluminium  acetate  is  used  as  a  mordant 


ALUMINIUM  571 

for  cotton  with  certain  dyes;  it  is  the  salt  of  a  weak  acid  and  is 
highly  hydrolyzed,  and  the  volatile  acid  formed  is  lost  by  evapora- 
tion when  the  cloth  dries. 

Aluminium  sulphate  is  used  in  making  white  leather.  The 
hide,  after  the  removal  of  the  hair,  is  soaked  in  a  bath  of  aluminium 
sulphate  and  salt.  The  change  that  takes  place  is  similar  to  that 
in  the  case  of  mordanting  wool.  The  animal  proteins  of  which  the 
skin  is  composed  absorb  aluminium  hydroxide  and  are  changed 
into  a  form  which  does  not  swell  in  water — the  skin  is,  thus, 
tanned  and  converted  into  leather.  Chromium  sulphate  is  used 
in  tanning  hides,  but  as  its  hydroxide  is  colored  it  cannot  be 
employed  when  white  leather  is  desired. 

Aluminium  salts  are  used  in  fire-proofing  and  water-proofing 
fabrics.  Either  is  accomplished  by  treating  the  fabric  with  alu- 
minium sulphate  and  then  with  a  solution  of  a  carbonate,  or  by 
wetting  it  with  a  solution  of  aluminium  acetate,  and  then  exposing 
it  to  steam;  the  acetate  is  hydrolyzed,  the  aluminium  hydroxide 
is  deposited  in  the  fiber,  and  the  acid  formed  passes  off  with  the 
steam. 

Alum,  K2SC>4,Al2(SO4)3,24H2O,  is  a  member  of  a  class  of 
double  sulphates  which  are  composed  of  1  molecule  of  a  sulphate 
of  an  alkali  metal  or  ammonium,  1  molecule  of  a  sulphate  of  a 
trivalent  metal,  (Al,Cr,Fe,Mn),  and  24  molecules  of  water  of 
crystallization;  all  the  members  of  this  series  of  compounds  are 
called  alums;  for  example,  ferric  ammonium  alum  has  the  formula 
(NH4)2SO4,Fe2(SO4)3,24H2O.  The  salts  crystallize  as  octahedra 
and  are  isomorphous. 

Alum  was  formerly  used  for  the  purposes  enumerated  above  in 
the  case  of  aluminium  sulphate,  because  it  crystallizes  well  from 
water  and  can  be  readily  obtained  in  a  pure  condition.  It  is  still 
used  to  some  extent,  but  improved  methods  of  preparation  of  the 
simple  sulphate  in  the  pure  condition  have  done  away  with  the 
necessity  of  using  the  more  expensive  salt  containing  potassium. 

689.  Clay. — When  feldspar  undergoes  decomposition  as  the 
result  of  atmospheric  influences,  it  is  converted  into  a  hydrated 
silicate  of  aluminium,  which  is  called  kaolin,  if  pure,  or  clay,  if  it  is 
mixed  with  other  substances.  The  relation  between  the  compo- 
sition of  feldspar  and  that  of  kaolin  can  be  seen  from  the  following 
formulas : 


572  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Feldspar,  KAlSi308  or  K20,Al203,6SiO2 

Kaolin    H2Al2(Si04)2,H20   or  2H2O,  Al203,2Si02 

The  potassium  silicate  produced  in  the  weathering  of  the  mineral 
is  soluble,  and  when  washed  away  leaves  pure  kaolin.  It  is  in  this 
way  that  soluble  potassium  salts  are  formed  in  the  soil.  If  feld- 
spar is  a  constituent  of  a  rock,  the  other  materials  remain  mixed 
with  the  clay  and  constitute  the  soil;  in  the  case  of  granite,  the 
silicon  dioxide  present  is  left  as  sand  and  the  mica  is  converted  into 
aluminium  silicate  and  potassium  silicate.  The  mineral  con- 
stituents of  soils  vary  with  the  composition  of  the  rock  from  which 
they  were  formed  and  the  extent  to  which  disintegration  of  the 
latter  has  taken  place.  The  presence  of  nitrates,  fluorides,  and 
phosphates,  and  of  calcium,  magnesium,  and  sodium  in  soil  has 
already  been  noted.  In  addition  to  these  mineral  constituents, 
soils  contain  a  mixture  of  complex  organic  substances,  known  as 
humus,  formed  as  a  result  of  the  disintegration  of  the  vegetable 
material  left  in  the  soil  when  plants  growing  on  it  die. 

Clays  may  remain  in  the  place  where  they  were  formed,  or  the 
kaolin,  which  is  light,  may  be  washed  away  by  streams  and  be 
deposited  in  another  place.  Such  sedimentary  clays  are  usually 
impure  as  the  result  of  the  admixture  of  other  substances;  they 
are  called  clays  if  sufficient  kaolin  is  present  in  them  to  form  a 
plastic  mass  when  they  are  mixed  with  water.  Impure  clays  of 
this  kind  are  used  in  making  brick,  sewer  pipe,  tile,  and  the  cheaper 
varieties  of  stoneware.  The  clay  is  made  into  a  plastic  mass  with 
water,  molded,  dried,  and  fired  in  a  kiln  at  about  1000°-1200°.  The 
substances  present  in  clay  other  than  kaolin  have  a  marked  effect 
on  the  properties  of  the  finished  product.  If  salts  of  sodium, 
potassium,  magnesium,  calcium,  or  iron  are  present,  they  react 
to  some  extent  at  the  temperature  of  "  burning  "  and  form  with 
the  kaolin,  silicates  which  fuse,  more  or  less.  In  this  way  they  bind 
together  the  clay  particles;  as  a  result,  the  material  after  "  burn- 
ing "  is  partially  vitrified  and  has  great  strength.  If  the  propor- 
tion of  the  metallic  compounds  is  small,  the  finished  product  is 
porous  and  lacks  luster. 

The  color  of  clay  products  is  due  to  the  presence  of  iron,  which 
is  present  in  the  clay,  in  all  probability,  as  a  colorless  hydrated 
silicate.  When  the  clay  is  "burned"  the  salt  is  decomposed,  and 


ALUMINIUM  573 

the  iron  is  oxidized  to  ferric  oxide,  Fe20s,  which  produces  a  red 
color.  If  the  clay  is  burned  in  a  reducing  atmosphere  ferrous 
silicate  is  formed  and  the  color  is  purplish  to  black.  If  iron  is 
absent,  the  clay  burns  white,  and  if  present  in  small  quantities, 
a  buff  color  is  produced. 

The  white  efflorescence  often  seen  on  brick  is  produced  as  the 
result  of  the  solvent  action  of  water  on  the  soluble  impurities  present 
in  the  clay,  which  have  not  been  converted  into  silicates  when  the 
clay  was  burned.  The  deposit  may  contain  the  sulphates  of 
sodium,  potassium,  magnesium,  and  calcium.  The  clay  products 
of  the  cheaper  varieties  are  glazed  by  introducing  salt  in  the  kiln 
after  the  material  has  been  fired;  at  the  high  temperature,  1200°, 
the  kaolin  on  the  surface  reacts  with  the  salt,  liberates  hydro- 
chloric acid,  and  is  converted  into  a  sodium-aluminium  silicate, 
which  fuses. 

690.  Cast  iron  is  covered  with  a  clay  enamel  in  making  bath- 
tubs, etc.,  by  coating  the  casting  with  a  mixture  prepared  from 
clay,  feldspar,  sand,  and  fluorspar  or  calcite;  cryolite  is  added  to 
render  the  enamel  opaque,  and  borax  makes  it  more  ductile  and 
elastic.     The  materials  are  first  melted  together  and  then  poured 
into  water  while  in  the  molten  condition.     This  treatment  con- 
verts the  product  into  a  finely  divided  condition;  it  is  then  mixed 
with  more  clay  and  ground  with  a  small  amount  of  water  until  it 
takes  the  form  of  a  thick  cream.     The  iron  casting  is  covered 
with  this  material  and  fired,  the  process  being  repeated  two  or 
more  times. 

691.  China  and  porcelain  are  prepared  from  pure  clay,  free 
from  iron,  to  which  has  been  added  more  or  less  feldspar,  the 
amount  being  determined  by  the  quality  of  ware  desired.     It  is 
not  necessary  in  burning  the  ware  containing  the  larger  propor- 
tion of  the  more  fusible  feldspar  to  heat  it  to  such  a  high  tem- 
perature, but  it  is  not  of  such  high  quality.     To  the  mixture  of 
finely  ground  material  is  added  just  enough  water  to  convert 
it  into  a  plastic  dough-like  mass,  which  is  then  shaped  either  on 
a  potter's  wheel  or  in  a  mold  made  of  plaster  of  Paris.     The 
material  is  next  allowed  to  dry  and  is  then  fired  in  a  kiln.     It 
is  glazed  by  dipping  it  into  a  cream  made  of  feldspar  and  silica, 
and  then  refiring  it  to  melt  the  feldspar,  which  flows  into  the 
pores  of  the  unglazed  material  and  produces  a  smooth  surface. 


574  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Lead  oxide  is  used  in  some  glazes  to  lower  the  temperature  of 
firing,  but  the  product  is  not  so  durable;  it  is  not  used  in  making 
the  highest  quality  china. 

692.  The   plasticity   of   clay — the   property   which   makes   it 
possible  to  convert  it  into  a  dough-like  mass  with  water — is  due 
to  the  colloidal  nature  of  the  particles  of  which  it  is  composed. 
These  are  in  a  very  finely  divided  condition  and  remain  more  or 
less  in  suspension  when  mixed  with  water,  on  account  of  the  fact 
that  they  are  charged  with  negative  electricity  and  repel  one 
another.     They  do  not,  therefore,  tend  to  coalesce  to  larger  par- 
ticles, which  would  settle  out  on  account  of  their  weight.     Clay  is 
rendered  colloidal  by  the  presence  of  certain  organic  substances 
produced  as  the  result  of  the  disintegration  of  the  organic  matter 
derived  from  plants.     This  material  often  gets  into  streams,  which 
are  thereby  rendered  muddy.     When  these  streams  enter  the  sea 
the  salts  present  cause  the  precipitation  of  the  colloidal  clay, 
which  often  forms  extensive  bars  at  the  mouths  of  rivers.     These 
facts  led  to  an  important  invention  by  Acheson,  who  first  pre- 
pared artificial  graphite  in  the  electric  furnace.     He  showed  that 
if  the  graphite  is  ground  with  tannin  or  other  organic  compounds 
similar  to  those  obtained  from  plants  which  render  clay  colloidal 
(534),  the  carbon  is   converted  into  a  colloidal  condition,   and 
when  mixed  with  water  or  oil  makes  an  excellent  lubricant  for 
many  purposes. 

Colloidal  clay  is  removed  from  water  by  adding  to  it  aluminium 
salts,  which  cause  the  precipitation  of  the  colloid.  A  process  is 
now  in  use  for  the  removal  of  ferric  oxide  from  clay,  which  is  based 
on  the  colloidal  properties  of  these  materials.  Clay  is  a  negative 
colloid,  being  charged  with  negative  electricity,  and  ferric  oxide  is  a 
positive  colloid.  When  an  electric  current  is  passed  through  a 
suspension  of  the  mixture,  the  clay  travels  with  the  negative  current 
and  accumulates  at  the  positive  pole,  and  the  ferric  oxide  moves  in 
the  opposite  direction.  The  material  which  settles  out  at  the 
positive  pole  is  free  from  iron  and  is  used  in  making  white  por- 
celain. 

693.  Fuller's  earth  is  a  form  of  clay  which  is  very  porous;  it  is 
used  as  an  absorbent  for  colored  substances  in  decolorizing  oils. 
Highly  colored  clays  are  used  as  pigments  in  making  paints; 
yellow  ocher  is  a  clay  colored  with  hydrated  oxide  of  iron,  and 


ALUMINIUM  575 

siennas  and  umbers  contain  in  addition  some  manganese  dioxide. 
If  the  material  is  heated  to  a  high  temperature  the  shade  is  deep- 
ened; red  ocher  is  made  in  this  way;  burnt  sienna  is  a  reddish 
orange. 

694.  Portland  Cement. — The  manufacture  of  cement  has 
become  one  of  the  most  important  chemical  industries,  on  account 
of  the  durable  properties  of  materials  constructed  from  cement  and 
the  fact  that  the  substances  from  which  it  is  prepared  are  abundant. 
The  Romans  used  for  building  purposes  a  rock  found  near  Mount 
Vesuvius,  which  had  the  property  of  hardening  under  water.  It 
was  a  rock  of  volcanic  origin,  which  was  a  natural  cement.  In 
1756  Smeaton  conducted  experiments  to  determine  the  best  kind 
of  mortar  to  be  used  under  water  in  connection  with  the  erection 
of  a  lighthouse  on  Eddystone  Rock,  and  found  that  one  made 
by  burning  lime  containing  clay  gave  the  best  results.  In  1824  a 
patent  was  granted  in  England  to  cover  the  manufacture  of  a 
cement  made  from  chalk  and  mud  rich  in  clay,  which  had  the 
property  of  setting  under  water;  it  was  called  Portland  cement 
because  when  it  set,  a  product  was  produced  which  resembled  in 
appearance  a  limestone  quarried  near  Portland,  England. 

Cement  is  formed  by  heating  together  substances  which  con- 
tain the  oxides  of  calcium,  aluminium,  and  silicon.  The  materials 
ordinarily  used  as  a  source  of  calcium  are  limestone,  chalk,  or 
marl,  which  is  an  amorphous  form  of  calcium  carbonate  mixed 
with  organic  matter  and  water.  The  aluminium  and  silicon  are 
obtained  from  clay  or  slate,  but  in  certain  cases  blast-furnace  slag 
is  now  used  in  large  quantities. 

Cement  consists  of  a  mixture  of  calcium  silicates  and  calcium 
aluminates  in  the  proportions  which  have  been  found  as  the  result 
of  practice  to  give  the  best  results.  The  materials  are  mixed  in 
such  proportions  that  the  percentage  of  CaO  in  the  mixture  is 
about  twice  the  sum  of  the  percentages  of  Si(>2  and  A^Os.  The 
relation  between  the  quantities  of  the  acid-forming  oxides  is 
approximately  1  AfeOa  to  2.5-4  SiCb. 

The  separate  materials  to  be  used  are  crushed,  dried,  and 
analyzed,  and  the  amounts  of  each  required  to  produce  a  cement 
of  the  above  composition  are  mixed.  The  charge  is  then  ground 
and  passed  through  a  rotary  kiln  from  70  to  150  feet  long  and  6  to  8 
feet  in  diameter  constructed  of  steel  and  lined  with  fire-brick. 


576  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  kiln  is  set  at  an  angle  of  about  15°  and  rotated  about  once  a 
minute.  As  the  powdered  material  tumbles  through  the  kiln  it  is 
heated  by  a  flame  produced  within  the  kiln  by  burning  gas,  oil,  or 
a  jet  of  powdered  coal.  The  material  is  heated  to  incipient  fusion 
and  forms  lumps.  The  "  clinker  "  so  produced  is  mixed  with 
about  2  per  cent  of  gypsum  and  ground  to  a  fine  powder.  The 
sulphate  is  added  to  retard  the  setting  of  the  cement  when  it  is 
mixed  with  water. 

The  chemical  composition  of  cement  and  the  changes  which  take 
place  when  it  sets  with  water  have  been  very  fully  studied,  but  no 
universally  accepted  explanation  of  these  changes  has  been  reached. 
It  is  thought  that  at  least  two  calcium  silicates  and  two  calcium 
aluminates  are  present  in  the  cement  before  it  is  treated  with 
water;  to  these  have  been  assigned  the  following  formulas: 
(CaO)3,Si02,  (CaO)2,Si02,  (CaO)3,Al203  and  (CaO)2,Al2O3.  The 
setting  of  cement  is  due  to  the  changes  which  take  place  when 
water  converts  one  or  more  of  these  substances  into  hydrates  and 
brings  about  their  hydrolysis  »with  the  liberation  of  free  calcium 
hydroxide. 

In  several  regions  in  the  United  States  rocks  occur  which  con- 
tain their  constituents  in  such  proportions  that  they  can  be  con- 
verted by  heating  directly  into  cement.  Such  rocks  occur  in 
Ohio,  Pennsylvania,  Illinois,  Wisconsin,  and  Colorado.  Large 
amounts  of  cement  are  made  from  blast-furnace  slag  (739)  and 
limestone. 

695.  Tests  for  Aluminium  Salts. — Solutions  of  alkalies,  car- 
bonates, and  sulphides  precipitate  from  salts  of  the  metal  color- 
less, gelatinous  aluminium  hydroxide,  which  is  soluble  in  sodium 
hydroxide  and  insoluble  in  ammonia.  When  barium  carbonate  is 
shaken  with  a  solution  of  a  salt  of  aluminium,  the  metal  is  pre- 
cipitated as  hydroxide  as  the  result  of  the  fact  that  the  hydrolysis 
of  the  salt  is  made  complete  owing  to  the  interaction  of  the  car- 
bonate with  the  acid  formed: 

A1C13  +  3H20  —  A1(OH)3  +  3HC1 
BaCO3  +  2HC1  ^±  BaCl2  +  H20  +  C02 

Salts  of  the  other  trivalent  metals  behave  in  a  similar  way,  but 
those  of  the  bivalent  metals  are  not  decomposed  because  they  are 


ALUMINIUM  577 

less  hydrolyzed  by  water.  The  reaction  serves,  therefore,  as  a 
means  of  separating  the  trivalent  and  bivalent  metals  and  is  used 
in  analytical  chemistry. 

EXERCISES 

1.  In  the  preparation  of  aluminium  by  electrolysis,  why  is  sodium  not  set 
free  from  the  electrolyte  used? 

2.  The  alloys  of  Al  and  Cu  vary  markedly  in  properties  as  the  proportions 
of  the  two  metals  vary.     This  is  not  the  case  with  the  Cu-Ag  alloys.     Can 
you  state  a  reason  for  the  difference? 

3.  What  compounds  could  be  used  in  place  of  the  oxides  of  iron  in  ther- 
mite?    Could  the  aluminium  be  replaced  by  any  other  metal?     If  so  what? 

4.  Starting  with  Al  write  equations  for  reactions  by  which  the  following 
compounds  could  be  prepared:    (a)  A1(NO3)3,    (6)  alum,    (c)  NaAlC>2,  and 
(d)  A1PO4,  which  is  insoluble. 

5.  Write  equations  for  the  reactions  which  take  place  between  the  follow- 
ing substances  in  aqueous  solution:    (a)  A1C13  and  Na2S,    (b)  A12(SO4)3  and 
Na2CO3,    (c)  A1(NO3)3  and  NaOH. 

6.  State  as  many  ways  as  possible  by  which  the  sulphates  of  the  following 
metals  can  be  distinguished  from  one  another:  Ca,  Mg,  Zn,  Al. 

7.  If  you  were  given  a  mixture  of  A1C13  and  MgCl2,  how  could  you  pre- 
pare from  it  pure  MgS04,  7H20  and  pure  A12(SO4)3,18H20? 


CHAPTER  XXXIX 
TIN  AND  LEAD 

696.  Tin  and  lead,  which  are  members  of  the  second  family  of 
the  fourth  group  in  the  periodic  classification  of  the  elements,  are 
characteristic  of  the  metals  which  show  the  valence  4.  The  com- 
pounds derived  from  them  in  the  bivalent  condition  resemble,  in 
general,  the  analogous  derivatives  of  other  metals  with  the  valence 
2,  but  then*  hydroxides  are  less  basic  than  those  of  the  metals  of 
the  calcium  family,  and  their  salts  are  more  highly  hydrolyzed  as 
a  consequence.  When  tin  and  lead  exhibit  the  valence  4  they 
function  as  acid-forming  elements;  stannic  chloride,  SnCU,  for 
example,  is  practically  not  ionized  in  solution,  and  is  slowly  but 
completely  hydrolyzed  by  water.  The  fact  that  increase  in 
valence  leads  to  an  increase  in  acid-forming  properties  has  been 
repeatedly  mentioned;  it  is  clearly  illustrated  in  the  case  of  these 
two  metals. 

Increase  in  atomic  weight  in  the  family  is  associated,  as  is  usual 
in  most  cases,  with  increase  in  base-forming  properties.  German- 
ium, the  first  member  of  the  family,  is  a  weak  acid-forming  element 
and  shows  more  resemblances  to  carbon  than  to  tin.  Its  monoxide, 
GeO,  is  practically  neutral  like  carbon  monoxide,  and  its  dioxide, 
GeO2,  is  acid-forming.  Lead  is  the  most  active  member  of  the 
family  in  base-forming  properties.  By  considering  the  above 
facts  in  connection  with  those  summarized  at  the  beginning  of 
the  last  chapter  in  regard  to  the  general  behavior  of  univalent, 
bivalent,  and  trivalent  metals,  it  will  be  seen  that  there  is  a  well- 
defined  gradation  in  chemical  properties  with  increase  in  valence. 
When  we  pass  to  the  next  group — the  fifth — the  members  of  which 
have  been  described,  we  find  that  the  basic  properties  of  the  ele- 
ments are  exceedingly  weak,  and  that  in  most  cases  the  important 
compounds  are  those  in  which  the  elements  function  as  acid- 
formers.  In  the  sixth  group,  which  contains  sulphur,  we  find  a 

578 


TIN  AND  LEAD  579 

strong  acid-forming  element,  but  some  of  the  elements  in  the  group 
in  their  lower  valencies  show  weakly  metallic  properties. 


TIN 

697.  Tin  was  known  to  the  ancients  and  was  used  in  bronze  in 
prehistoric  times.     The  metal  is  obtained  from  cassiterite,  which 
is  commonly  called  tin-stone,  SnO2.    The  ore  is  first  roasted  to 
remove  sulphur,  arsenic,  and  antimony,  and  to  oxidize  any  iron 
present.     It  is  then  treated  with  hydrochloric  acid  to  convert  the 
impurities  in  the  ore  into  chlorides,  which  are  washed  out  with 
water.     The  oxide  of  tin  is  then  reduced  with  coal.     Large  quanti- 
ties of  tin-stone  were  formerly  mined  at  Cornwall,  England,  but 
the  chief  sources  of  the  metal  at  present  are  the  Malay  Peninsula, 
Bolivia,  and  the  island  of  Banca  in  the  Indian  Ocean. 

The  chief  physical  and  chemical  properties  of  the  metal  have 
already  been  given  (table,  page  443  and  548).  It  exists  in  two 
forms.  One  is  amorphous,  has  the  specific  gravity  5.85,  and  is 
stable  below  18°;  the  other,  which  is  crystalline  and  stable  above 
18°,  can  exist  below  this  temperature,  but  when  it  is  kept  at  —15° 
it  passes  slowly  into  the  amorphous  form.  Tin  has  a  low  tenacity, 
but  is  very  malleable  and  can  be  made  into  foil  as  thin  as  0.01  of 
an  inch.  It  is  most  malleable  at  about  100°,  and  becomes  so 
brittle  at  200°  it  can  be  powdered.  It  melts  at  232°,  boils  at  2270° 
and  has  the  specific  gravity  7.3.  The  action  of  the  air  and  acids 
on  tin  has  been  described  (548,  550,  549). 

Tin  is  used  under  the  name  block  tin  or  sheet  tin  in  the  manu- 
facture of  pipe  and  of  vessels  of  various  kinds,  when  a  metal  is 
desired  which  is  not  attacked  by  dilute  vegetable  acids.  On 
account  of  the  high  cost  of  tin  other  metals,  such  as  copper  and 
iron,  are  coated  with  it.  Tin-plate  is  made  by  coating  sheets  of 
steel  with  molten  tin  (551) ;  it  is  apt  to  corrode  rapidly  when  the 
steel  is  exposed  in  any  way  (567) .  Large  quantities  of  tin  are  used 
in  making  alloys,  such  as  solder  and  bronze  (542) . 

698.  Oxides  and  Hydroxides  of  Tin. — When  stannous  salts 
are  treated  in  solution  with  bases,  stannous  hydroxide  is  formed  as  a 
colorless  precipitate: 

SnCl2  +  2NaOH  =  Sn(OH)2  +  2NaCl 


580  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  hydroxide  dissolves  in  an  excess  of  an  alkali  to  form  a  stannite : 
Sn(OH)2  +  2NaOH  =  Sn(ONa)2  +  2H2O 

The  salt  is  highly  hydrolyzed  and  is  unstable;  when  the  solution 
is  boiled,  it  is  converted  into  tin  and  sodium  stannate,  which 
resembles  sodium  carbonate  in  composition,  Na2SnO3,  and  is 
derived  from  SnO2 : 

2Sn(ONa)2  +  H2O  =  Sn  +  Na2SnO3  +  2NaOH 

Stannous  hydroxide  is  precipitated  when  a  soluble  carbonate 
is  added  to  a  stannous  salt,  the  reaction  being  produced  by  the 
same  cause  that  leads  to  the  precipitation  of  the  hydroxides  of 
the  trivalent  metals  by  carbonates  (682).  The  base-forming 
property  of  stannous  tin  is  exceedingly  weak.  The  fact  is  of 
interest  in  connection  with  the  behavior  of  salts  of  other  bivalent 
metals  with  soluble  carbonates;  those  of  calcium  yield  the  normal 
carbonate,  those  of  zinc,  a  basic  carbonate,  and  those  of  stannous 
tin,  the  base  itself. 

Stannous  oxide,  SnO,  is  obtained  by  heating  stannous  oxalate, 
SnC2O4,  in  the  absence  of  oxygen;  it  is  a  black  powder,  which 
burns  in  the  air  to  the  white  dioxide,  Sn02. 

699.  When  bases  are  added  to  a  solution  of  a  stannic  com- 
pound,  SnCU,   for  example,   a  white,   gelatinous  precipitate  is 
formed,  which  is  an  unstable  hydroxide  that  loses  water  gradually 
until  stannic  oxide,  SnO2,  is  formed.     The  precipitate  dissolves 
in  solutions  of  alkalies  and  yields  salts  of  stannic  acid,  which  is 
derived  from  the  normal  hydroxide  by  loss  of  water: 

Sn(OH)4  -  H2O  =  H2SnO3 

Stannic  acid  and  its  salts  are  sometimes  distinguished  from  a 
second  stannic  acid,  formed  by  the  action  of  concentrated  nitric 
acid  on  tin,  by  prefixing  to  their  names  the  Greek  letter  alpha,  a. 

700.  When  tin  is  treated  with  concentrated  nitric  acid  it  is 
converted  into  a  hydrated  derivative  of  the  dioxide,  which  differs 
in  properties  from  the  stannic  acid  precipitated  by  alkalies  from 
stannic  compounds;    it  is  called  metastannic  acid  or  /3-stannic 
acid.     It  dissolves  with  difficulty  in  boiling  solutions  of  alkalies, 
from  which  salts  can  be  obtained;   the  sodium  salt  made  in  this 
way  has  the  composition  represented  by  the  formula 


TIN  AND  LEAD  581 

For  this  reason  the  formula  of  metastannic  acid  is  generally 
written  (SnO2)5,^*H20;  the  amount  of  water  in  combination  varies 
with  the  conditions  under  which  the  acid  is  made.  When  the 
acid  is  fused  with  sodium  hydroxide,  it  is  converted  into  sodium 
stannate,  Na2SnOs,  and  when  heated  alone,  into  the  dioxide. 
Tin  dioxide  is  white  when  cold,  but  resembles  zinc  oxide  in  turning 
yellow  when  hot.  The  oxide  that  occurs  as  cassiterite  is  black, 
the  color  being  due  to  the  presence  of  oxides  of  iron. 

Tin  dioxide  is  used  in  fire-proofing  fabrics.  The  materials 
are  first  soaked  in  a  solution  of  sodium  stannate  and  then  treated 
with  a  solution  of  ammonium  sulphate.  The  reaction  which  takes 
place  leads  to  the  precipitation  in  the  fiber  of  stannic  acid,  which 
loses  water  and  passes  into  the  dioxide. 

701.  The  Chlorides  of  Tin. — Stannous  chloride  is  formed  when 
tin  dissolves  in  hydrochloric  acid.  The  salt  obtained  from  the 
solution  by  crystallization  has  the  formula  SnCl2,2H2O.  It  is 
highly  hydrolyzed,  and  when  a  strong  solution  of  it  is  poured  into 
water  a  precipitate  of  a  basic  salt,  Sn(OH)Cl,  is  formed. 

Stannous  chloride  is  oxidized  by  the  air  and  is  slowly  converted 
in  solution  into  stannic  chloride,  SnCU,  and  basic  stannous  chloride, 
Sn(OH)Cl,  which  precipitates.  The  salt  is  kept  in  the  stannous 
condition  and  the  precipitation  of  the  basic  salt  avoided  by  adding 
to  the  solution  a  small  amount  of  hydrochloric  acid  and  some 
metallic  tin 

The  tendency  of  stannous  tin  to  pass  into  stannic  tin  is  also 
shown  when  it  is  brought  into  contact  with  salts  that  can  be 
reduced.  Ferric  chloride,  for  example,  is  reduced  to  ferrous  chlo- 
ride—2FeCl3  +  SnCl2  =  2FeCl2  +  SnCU— and  mercuric  chloride, 
to  mercurous  chloride  and  mercury  (674).  Stannous  chloride  is 
used  as  a  mordant  under  the  name  tin  crystals. 

Stannic  chloride  is  made  from  tin  and  chlorine.  The  anhy- 
drous compound  is  a  colorless,  fuming  liquid,  which  boils  at  114° 
and  dissolves  in  water.  It  forms  a  number  of  hydrates,  of  which 
the  one  having  the  formula  SnCl4,5H20  is  the  most  important. 
It  is  used  as  a  mordant  on  account  of  the  fact  that  it  is  appreciably 
hydrolyzed  and  the  hydroxide  formed  is  absorbed  by  the  fiber  and 
holds  the  dye.  Large  quantities  of  it  are  used  in  making  the 
cheaper  grades  of  silk,  because  it  adds  materially  to  the  weight 
of  the  latter  and  improves  its  appearance.  Silks  weighted  in  this 


582  INORGANIC  CHEMISTRY  FOR  COLLEGES 

way  contain  at  times  as  much  as  75  per  cent  of  the  oxide  and  are 
not  durable.  "  Pink  salt,"  which  is  extensively  used  as  a  mordant, 
is  a  double  salt  of  the  formula  (NH^SnCle.  A  large  number  of 
double  salts  containing  stannic  tin  are  known;  their  easy  forma- 
tion is  traceable  to  the  acid-forming  properties  of  the  metal. 

702.  The  Sulphides  of  Tin. — Stannous  sulphide,  SnS,  obtained 
by  the  action  of  hydrogen  sulphide  on  a  solution  of  stannous 
chloride1  is  a  brown  precipitate,  which  is  insoluble  in  ammonium 
sulphide  but  dissolves  in  ammonium  polysulphide  as  the  result 
of    the    formation    of    ammonium    thiostannate,     (NH^SnSa. 
Addition  of  an  acid  causes  the  precipitation  of  yellow  stannic 
sulphide,  SnS2,  which  is  also  formed  from  stannic  chloride  and 
hydrogen  sulphide.     The  reactions  are  analogous  to  those  dis- 
cussed at  length  in  the  case  of  arsenic  and  antimony  (478)  and, 
as  a  result,  in  qualitative  analysis  tin  falls  into  the  group  which 
includes  these  elements. 

703.  Test  for  Tin  Salts. — The  properties  of  the  sulphides  just 
given  are  utilized  in  analytical  chemistry.     A  confirmatory  test 
lor  tin  is  based  on  the  fact  that  metallic  zinc  sets  free  tin  from  its 
salts,  and  the  metal  obtained  in  this  way  is  soluble  in  dilute  hydro- 
chloric acid. 

LEAD 

704.  Metallurgy  of  Lead. — Lead  has  been  known  since  pre- 
historic times.    It  occurs  as  the  carbonate  and  the  sulphate,  which 
are  used  as  ores  of  the  metal,  but  the  most  important  source  is 
the  sulphide,  PbS,  galena,  which  crystallizes  in  black  cubes  having 
a  high  silvery  luster.     The  metal  is  ordinarily  obtained  from  the 
ore  by  first  heating  it  in  a  stream  of  air  to  convert  it  into  the 
oxide,  and  then  reducing  the  latter  in  a  blast  furnace  with  coal. 
Many  ores  of  lead  contain  silver  and  gold,  and  a  part  of  the 
metallurgical  processes  used  in  extracting  lead  have  to  do  with  the 
separation  of  the  noble  metals;    these  will  be  considered  later 
(713). 

705.  Properties  of  Lead. — The  more  important  physical  prop- 
erties of  lead  have  already  been  given  (table,  page  443).     Lead 
has  a  bluish-gray  color  with  a  marked  metallic  luster  when  freshly 
cut,  but  it  soon  tarnishes  in  the  air  as  the  result  of  the  formation 
of  a  thin  coating  of  oxide,  which  changes  to  a  basic  carbonate. 


TIN  AND  LEAD  583 

The  coating  is  closely  adherent  and  retards  further  corrosion  to 
such  an  extent  that  lead  is  used  when  a  metal  is  required  that  is 
stable  under  atmospheric  influences.  The  metal  is  very  inactive 
toward  acids,  and  is  used  in  the  manufacture  of  many  kinds  of 
apparatus  designed  for  the  chemical  industries.  Its  use  in  alloys 
has  already  been  explained  (542).  Lead  is  used  in  large  quantities 
in  the  manufacture  of  pipe,  which  is  made  by  forcing  the  metal 
while  soft  through  dies. 

706.  Oxides  of  Lead. — When  the  metal  is  heated  in  the  air  it  is 
converted  at  a  temperature  just  above  its  melting-point,  327°, 
into  a  yellow  monoxide,  PbO,  which  is  called  massicot.  If  the 
temperature  is  such  that  the  oxide  melts,  the  product  contains  a 
trace  of 'red  lead  and  is  reddish  yellow;  it  is  known  as  litharge. 

Red  lead,  PbsO4,  is  prepared  by  first  converting  lead  into  massi- 
cot and  then  heating  it  at  about  480°  in  air,  the  process  requiring 
about  forty-eight  hours.  This  oxide  is  used  in  glass  making  and 
in  paints  to  protect  iron  and  steel  from  rusting.  It  reacts  with 
linseed  oil  to  form  a  hard  solid,  and  for  this  reason  a  mixture 
of  the  two  substances  is  used  as  a  lute  in  plumbing  and  gas  fitting. 
A  mixture  of  red  lead  and  glycerine  is  used  as  a  cement,  because  on 
standing  it  is  converted  into  a  hard,  adhering  mass. 

When  red  lead  is  heated  to  a  high  temperature  it  gives  off 
oxygen  and  is  converted  into  the  monoxide.  Warm  dilute  nitric 
acid  converts  it  into  lead  nitrate,  Pb(NOs)2,  which  dissolves,  and 
lead  dioxide,  PbO2,  which  is  left  as  a  brown  powder.  For  this 
reason  red  lead  is  thought  to  be  a  salt  of  the  formula  Pb2(Pb04), 
which  is  derived  from  a  hydroxide  of  the  composition  Pb(OH)4  by 
replacing  the  hydrogen  atoms  by  two  atoms  of  bivalent  lead. 
When  lead  has  the  valence  2  it  has  base-forming  properties;  in 
the  quadrivalent  condition  it  is  acidic  and  its  hydroxide,  Pb(OH)4, 
is  called  orthoplumbic  acid.  From  this  point  of  view  the  reaction 
between  nitric  acid  and  red  lead,  which  is  lead  orthoplumbate,  is 
written  as  follows : 

Pb2Pb04  +  4HN03  =  2Pb(N03)2  +  H4Pb04 
H4PbO4  =  Pb02  +  2H2O 

Lead  dioxide ,  Pb02,  is  a  brown  powder,  insoluble  in  water, 
which  is  usually  prepared  by  treating  with  bleaching  powder  a 


584  INORGANIC  CHEMISTRY  FOR  COLLEGES 

solution  of  sodium  plumbite,  Na2pb02,  formed  by  dissolving  lead 
hydroxide  in  caustic  soda.  The  sodium  salt  hydrolyzes  giving 
sodium  hydroxide  and  lead  hydroxide,  which  is  oxidized  by  the 
oxygen  furnished  by  the  bleaching  powder :  Na2PbO2  +  CaOCl2  + 
H2O  =  2NaOH  +  CaCl2  +  PbO2.  The  dioxide  dissolves  in  strong 
solutions  of  alkalies,  playing  the  part  of  an  acid  anhydride,  and 
forms  salts  related  in  composition  to  the  stannates  and  carbonates. 
Potassium  plumbate  is  obtained  from  aqueous  solutions  in  crystals, 
which  have  the  formula  K2PbO3,3H2O. 

Lead  dioxide  reacts  with  hydrochloric  acid  in  a  way  entirely 
analogous  to  that  which  has  been  described  in  the  case  of  man- 
ganese dioxide  (112):  PbO2  +  4HC1  =  PbCl2  +  2H2O  +  C12. 
The  use  of  lead  dioxide  in  storage  batteries  has  been  described  at 
some  length  in  section  578. 

Lead  hydroxide,  Pb(OH)2,  is  formed  as  a  white  precipitate 
when  a  solution  of  a  base  is  added  to  a  lead  salt;  it  reacts  with 
caustic  alkalies  and  forms  soluble  plumbites,  but  does  not  dissolve 
in  ammonia. 

707.  The  Chlorides  of  Lead.— Lead  chloride,  PbCl2,  is  diffi- 
cultly soluble  in  cold  water  and  is  precipitated  when  a  soluble 
chloride  is  added  to  a  solution  of  a  lead  salt ;  it  is  colorless  and  dis- 
solves in  hot  water,  from  which  it  crystallizes  when  the  solution 
cools. 

Lead  tetrachloride,  PbCU,  is  stable  only  at  low  temperatures; 
it  is  a  liquid  which  fumes  in  the  air  and  decomposes  with  a  small 
amount  of  water  to  form  the  dichloride  and  chlorine.  With  larger 
amounts  of  water  it  is  hydrolyzed  and  converted  into  lead  dioxide 
and  hydrochloric  acid.  Double  salts  containing  lead  tetrachloride 
are  comparatively  stable,  and  advantage  is  taken  of  this  fact  in 
preparing  the  compound.  When  chlorine  is  passed  into  a  mixture 
of  lead  chloride  and  hydrochloric  acid,  the  lead  tetrachloride 
formed  unites  with  some  of  the  acid  and  forms  a  complex  acid  of 
the  composition  H2PbCle  (PbCU,2HCl).  Ammonium  chloride  is 
next  added  to  the  solution  to  form  ammonium  chloroplumbate, 
(NH4)2PbCl6.  When  this  salt  is  added  to  concentrated  sulphuric 
acid  at  a  low  temperature,  the  salt  is  decomposed  and  lead  tetra- 
chloride separates  as  a  heavy  oil. 

Lead  iodide,  PbI2,  has  characteristic  physical  properties  and 
is  useful  in  identifying  lead  compounds.  It  is  precipitated  as  a 


TIN  AND  LEAD  585 

yellow  powder  when  a  soluble  iodide  is  added  to  a  lead  salt.  It 
crystallizes  from  boiling  water  in  yellow  scales,  which  have  a  brilliant 
luster. 

708.  Lead  Carbonate. — The  neutral  carbonate  of  lead,  PbCOs, 
occurs  as  the  mineral  cerussite,  and  was  used  by  the  Romans  as  a 
white  pigment.  It  was  early  replaced  by  a  basic  carbonate  for 
this  purpose,  and  a  method  of  preparing  the  latter  used  in  Hol- 
land in  the  seventeenth  century  is  still  employed  in  the  manu- 
facture of  "  white  lead,"  which  is  to-day  the  most  important 
pigment  used  in  making  paint.  The  value  of  a  pigment  is  largely 
determined  by  its  so-called  " covering  power,"  which  depends 
on  its  opacity  and  physical  condition.  White  lead  as  ordinarily 
made  has  a  composition  closely  approaching  that  of  the  formula: 
Pb3(OH)2(CO3)2  or  Pb(OH)2,2PbCO3.  A  compound  of  the 
same  formula  is  obtained  when  sodium  carbonate  is  added  to  a 
soluble  lead  salt,  but  it  does  not  have  the  same  covering  power  as 
white  lead.  Many  attempts  have  been  made  to  replace  the  old 
Dutch  process  for  making  the  pigment,  for  it  requires  a  long  time. 
Some  of  these  have  been  successful,  but  the  white  lead  prepared 
by  the  older  process  is  held  by  some  to  have  superior  qualities. 

In  the  Dutch  process  for  making  white  lead,  sheet  lead  or  cast- 
ings made  from  the  metal  in  such  a  form  that  they  have  a  large 
surface,  are  placed  on  shelves  in  earthenware  pots,  the  bottoms  of 
which  are  covered  with  a  dilute  solution  of  acetic  acid,  HC2H3O2. 
The  pots  are  set  on  a  layer  of  moist  spent  tan  bark.  A  floor  is 
next  placed  a  few  inches  above  the  tops  of  the  open  pots,  and  this 
is  covered  with  tan  bark,  upon  which  more  pots  are  set.  The 
process  is  repeated  until  the  room  is  filled.  It  is  then  closed  and 
left  for  about  three  months.  The  bark  soon  begins  to  ferment 
and,  as  the  result,  carbon  dioxide  is  formed  and  the  temperature 
rises  to  55°-60°.  Loss  of  heat  is  prevented  by  having  a  thick 
layer  of  the  bark  around  the  walls  of  the  room.  Acetic  acid  is 
volatilized  and  slowly  converts  the  lead  into  an  acetate,  which  is 
changed  by  the  water  and  carbon  dioxide  into  the  basic  carbonate, 
the  acetic  acid  being  regenerated.  When  the  corroded  material 
is  taken  from  the  pots  it  is  found  to  be  in  the  form  of  a  hard  com- 
pact mass,  which  retains  the  original  shape  of  the  lead.  It  is 
ground  to  a  fine  powder  under  water,  and  the  milky  liquid  is  allowed 
to  settle.  After  the  larger  particles  have  separated,  the  liquid  is 


586  INORGANIC  CHEMISTRY  FOR  COLLEGES 

withdrawn,  and  on  further  settling  the  white  lead  is  obtained  as  a 
heavy  mud.  This  is  then  mixed  with  a  little  linseed  oil,  which 
causes  the  separation  of  the  water.  The  mixture  of  white  lead 
and  oil  is  separated  and  packed.  This  method  of  treating  the 
mud  avoids  handling  dry,  powdered  material,  which  is  poisonous 
and  is  apt  to  get  into  the  air  as  dust. 

In  Carter's  process,  which  is  much  used,  lead  is  first  converted 
into  a  very  finely  divided  condition  by  blowing  a  jet  of  super- 
heated steam  against  a  stream  of  molten  lead.  The  "  atomized  " 
lead  is  then  tumbled  in  a  rotating  cylinder  into  which  acetic  acid 
and  carbon  dioxide  are  led.  The  corrosion  is  complete  in  about 
fifteen  days. 

White  lead  is  an  excellent  pigment  and  is  the  basis  for  most 
colored  paints,  which  are  prepared  by  the  addition  of  colored 
substances  to  a  mixture  of  white  lead  and  linseed  oil.  White 
lead  is  converted  by  hydrogen  sulphide  into  lead  sulphide,  which  is 
black;  it  cannot  be  used  in  paints  which  are  exposed  to  this  gas. 
Under  these  circumstances  zinc  white,  which  forms  a  white  sul- 
phide, is  used. 

709.  Other  Salts  of  Lead.— Lead  nitrate,  Pb(NO3)2,   is  the 
common  laboratory  reagent.     It  is  made  by  dissolving  lead  oxide . 
in  nitric  acid;    it  crystallizes  readily  from  hot  water  containing 
a  small  amount  of  nitric  acid,  which  prevents  its  hydrolysis. 

Lead  sulphate,  PbSO*,  is  a  very  insoluble  salt  and  is  made  by 
precipitation;  it  dissolves,  like  barium  sulphate,  in  concentrated 
sulphuric  acid.  It  is  soluble  in  concentrated  solutions  of  the 
alkalies  as  the  result  of  the  formation  of  sodium  plumbite. 

Lead  acetate,  Pb^HsC^,  called  sugar  of  lead,  is  prepared 
by  the  action  of  air  and  acetic  acid  on  lead,  or  by  dissolving 
lead  oxide  in  the  acid.  It  is  used  in  the  preparation  of  a  basic 
lead  acetate,  which  is  formed  when  a  solution  of  the  neutral 
salt  is  boiled  with  litharge.  Basic  lead  acetate  has  the  formula 
Pb(OH)C2H3O2;  it  is  soluble  in  water  and  is  used  as  a  mordant. 

Lead  sulphide,  PbS,  is  formed  as  a  black  precipitate  from  lead 
salts  and  soluble  sulphides.  It  is  insoluble  in  dilute  hydrochloric 
acid,  but  dissolves  in  concentrated  nitric  acid. 

710.  Tests  for  Lead  Salts. — The  formation  of  a  black  sulphide 
in  the  presence  of  dilute  acids  when  a  lead  salt  is  treated  with 
hydrogen  sulphide,  and  the  formation  of  a  white  insoluble  sulphate 


TIN  AND  LEAD  587 

when  lead  salts  are  treated  with  a  soluble  sulphate,  are  used  as  a 
test  for  the  metal.  Sodium  hydroxide  precipitates  white  lead 
hydroxide  which  is  soluble  in  an  excess  of  the  reagent  but  insoluble 
in  ammonium  hydroxide.  Lead  appears  in  the  first  group  in  qual- 
itative analysis  as  it  is  precipitated  as  chloride  when  hydrochloric 
acid  is  added  to  a  solution  containing  it.  It  is  separated  from  the 
insoluble  chlorides  by  dissolving  it  from  the  mixture  with  boiling 
water,  from  which  it  separates  on  cooling.  Since  lead  chloride  is 
slightly  soluble  in  water  (1.5  parts  in  100  at  18°),  lead  is  not  com- 
pletely removed  from  solution  when  it  is  precipitated  as  chloride. 

EXERCISES 

1.  Summarize  in  the  form  of  a  table  the  properties  of  metallic  elements 
and  their  hydroxides  and  salts,  which  are  markedly  influenced  by  the  valence 
of  the  metal.     Consider  metals  having  the  valence  1,  2,  3,  and  4. 

2.  Write  equations  for  reactions  which  take  place  between  the  following: 
(a)  Sn  and  concentrated  HNO3,   (6)  Sn  and  dilute  HNO?,   (c)  SnCl2,  H2O,  and 
air,    (d)  Na2SnO3  and  (NH4)2SO4. 

3.  By  what   chemical  properties  could  you  distinguish    (a)  Sb2S3  from 
SnS2,    (6)  SnCl2  from  A1C13,    (c)  pure  tin  foil  from  a  foil  containing  Sn  and 
Pb,   (d)  pure  tin.  from  an  alloy  of  Sn  and  Zn? 

4.  How  could  you  determine   (a)  whether  a  sample  of  silk  had  been  mor- 
danted with  tin  salt  and   (6)  the  percentage  of  SnO2  present? 

5.  How  could  you  determine  whether  a  piece  of  cotton  had  been  fire- 
proofed  with  alum  or  sodium  stannate? 

6.  Write  equations  for  reactions  which  would  serve  to  distinguish  from 
one  another  the  following:   (a)  PbO,    (6)  PbO2,  and   (c)  Pb3O4. 

7.  How  could  you  obtain    (a)  Pb3O4  from  a  mixture  of  Pb3O4  and  ZnO, 
(6)  MgO  from  a  mixture  of  MgO  and  PbO? 

8.  How  could  Pb  and  Cu  be  separated? 

9.  (a)  Why  does  lead  form  double  salts  readily?     What  substance  could 
be  used  to  dissolve  lead  iodide? 

10.  (a)  State  a  number  of  ways  of  determining  whether  a  white  paint 
is  made  from  white  lead  or  zinc  white.     (6)  If  the  paint  contains  both,  how 
could  you  show  the  presence  of  each? 

11.  What  substance  could  be  used  to  restore  the  white  color  to  a  lead 
paint  which  has  turned  yellow  in  the  air? 


CHAPTER  XL 
COPPER,  SILVER,  AND  GOLD 

711.  The  position  of  copper,  silver,  and  gold  in  the  first  group 
of  the  periodic  classification  is  anomalous,  because  the  metals 
show  no  marked  resemblances  to  those  of  the  alkalies,  and  because 
copper  and  gold  form  compounds  in  which  the  metals  have  a 
greater  valence  than  1.       It  is  for  this  reason  that  their  consider- 
ation has  been  delayed  to  this  point.     We  have  seen  that  there  is  a 
gradual  change  in  the  properties  of  the  compounds  derived  from 
metals  as  the  valence  of  the  latter  changes  from  1  to  4;   and  we 
have  learned  the  general  properties  of  the  compounds  of  the  four 
types  of  metals.    With  these  facts  in  mind  it  will  be  easier  to  under- 
stand the  chemical  behavior  of  copper,  which  functions  as  a  uni- 
valent  and  bivalent  metal,  and  that  of  gold,  which  has  the  valencies 
1  and  3.     Silver  is  always  univalent  and  the  properties  of  its  com- 
pounds are  typical  of  those  of  other  compounds  derived  from  heavy 
metals  when  they  exhibit  this  valence.     For  this  reason  silver 
and  its  compounds  will  be  first  considered. 

SILVER 

712.  Silver  is  a  relatively  strong  base-forming  element;    its 
hydroxide,  which  exists  only  in  solution  and  is  formed  when  silver 
oxide  dissolves  in  water,  is  highly  dissociated  and  forms  salts 
which  are  not  hydrolyzed.     In  this  respect  it  resembles  the  alkali 
metals.     It  is  a  very  inactive  element,  however;  it  does  not  liberate 
hydrogen  from  acids,  and  its  oxide  is  readily  decomposed  by  heat. 

The  metal  has  been  known  since  the  earliest  times  and  was 
studied  by  the  alchemists,  who  used  a  crescent  as  its  symbol  in 
their  writings.  The  symbol  used  to-day,  Ag,  is  derived  from  the 
Latin  name  of  the  metal  (argentum). 

588 


COPPER,  SILVER,  AND  GOLD  589 

713.  Occurrence    and    Metallurgy    of    Silver. — The    metal 
occurs  in  the  free  condition  often  alloyed  with  gold  and  with 
copper.     It  occurs  as  the  sulphide,  Ag2S,  argentite,  which  is  asso- 
ciated with  lead  sulphide.     The  metal  is  obtained  in  large  quan- 
tities as  a  by-product  in  the  smelting  of  lead  and  copper.     It  also 
occurs  as  double  sulphides  with  antimony  and  arsenic,  and  as 
silver  chloride. 

Native  silver  is  obtained  by  amalgamation ;  the  crushed  ore  is 
allowed  to  stand  in  contact  with  mercury,  which  dissolves  the 
metal,  and  the  solution  formed  is  then  distilled;  the  residue  is 
silver  alloyed  with  any  gold  present  in  the  ore.  The  method  most 
commonly  used  when  the  ore  is  a  compound  of  the  metal,  is  to 
extract  it,  after  pulverization,  with  a  dilute  solution  of  sodium 
cyanide.  The  silver  is  converted  into  silver  cyanide,  which  dis- 
solves in  an  excess  of  the  alkali  cyanide  to  form  the  complex  salt, 
NaAg(CN)2.  Zinc  is  then  added  to  the  solution  to  precipitate  the 
metal.  In  the  case  of  some  ores,  the  cyanide  and  amalgamation 
processes  are  used  simultaneously. 

Low-grade  silver  ores  are  often  smelted  with  lead  ores,  since 
the  latter  frequently  contain  silver,  which  must  be  recovered.  The 
method  commonly  employed  is  called  the  Parkes  process.  Silver 
and  gold  are  much  more  soluble  in  zinc  than  in  lead;  and  since 
zinc  is  practically  insoluble  in  lead  it  can  be  used  to  extract  the 
precious  metals  from  the  latter.  In  carrying  out  the  process 
zinc  is  well  stirred  into  the  molten  lead,  which  is  then  allowed  to 
cool  slowly.  A  scum  consisting  of  zinc,  silver,  gold,  and  some  lead 
separates  on  the  surface  of  the  metal;  this  is  removed  and  dis- 
tilled and  thus  freed  from  zinc.  The  residue  is  then  "  cupelled," 
that  is,  it  is  heated  in  a  stream  of  air,  which  oxidizes  the  lead  and 
blows  the  molten  oxide  to  the  edge  of  the  hearth,  where  it  is 
skimmed  off. 

Processes  which  are  used  in  Mexico  and  South  America  are 
based  on  other  reactions.  In  one  of  these  the  ground  ore  is  mixed 
with  water,  salt,  and  copper  sulphate,  and  heated  for  several  hours. 
Metallic  iron  is  then  added  to  set  silver  free  from  the  chloride,  and 
the  metal  is  extracted  from  the  mass  by  amalgamation. 

714.  Properties  and  Uses  of  Silver. — The  chief  physical  prop- 
erties of  silver  are  given  in  the  table  on  page  443  and  discussed  in 
section  548.    The  metal  is  the  best  conductor  of  electricity  and 


590  INORGANIC  CHEMISTRY  FOR  COLLEGES 

heat;  it  is  exceedingly  malleable  and  ductile.  The  molten  metal 
possesses  the  property  of  dissolving  oxygen  (about  22  volumes), 
which  is  expelled  when  it  solidifies.  In  the  massive  or  crystalline 
condition  silver  is  white;  when  precipitated  from  solution  by  zinc 
it  forms  a  black  powder,  but  mercury  causes  the  separation  of  the 
metal  from  its  salts  in  a  highly  crystalline  condition  which  leads 
to  the  formation  of  what  is  called  a  silver  "  tree."  By  the  use  of 
reducing  agents  silver  can  be  obtained  in  the  colloidal  condition, 
the  color  of  the  metal  varying  with  the  substance  used. 

Silver  is  a  relatively  soft  metal  and  for  this  reason  an  alloy 
of  the  metal  is  used  in  coins.  Those  of  the  United  States  and  the 
countries  of  continental  Europe  contain  900  parts  of  silver  and  100 
parts  of  copper;  the  coins  of  Great  Britain  and  sterling  silver 
contain  925  parts  of  silver  and  75  parts  of  copper.  So-called 
oxidized  silver  is  made  by  treating  the  metal  with  a  solution  of 
a  soluble  sulphide,  which  causes  the  precipitation  of  silver  sul- 
phide on  the  surface  of  the  metal.  In  making  mirrors  silver  is 
deposited  on  glass  by  covering  the  well-cleaned  surface  with 
a  solution  of  a  silver  salt  to  which  has  been  added  a  reducing 
agent,  such  as  glucose  or  an  ammoniacal  solution  of  cream  of 
tartar. 

Silver  is  a  very  inactive  element  and  does  not  tarnish  in  the 
air  except  in  the  presence  of  hydrogen  sulphide.  It  turns  yellow 
or  black  when  in  contact  with  certain  organic  products  which 
contain  sulphur,  such  as  the  protein  in  eggs.  The  behavior  of  the 
metal  with  acids  has  been  discussed.  It  dissolves  in  nitric  and 
concentrated  sulphuric  acids;  it  is  not  attacked  by  alkalies,  even 
when  they  are  fused,  because  it  does  not  play  the  part  of  an  acid- 
forming  element. 

Large  quantities  of  silver  are  used  in  plating  other  metals  (583) 
and  in  the  preparation  of  silver  nitrate,  which  is  the  source  of  the 
compounds  of  the  metal  used  in  photography. 

715.  Oxides  of  Silver. — When  a  base  is  added  to  a  solution  of  a 
silver  salt  a  brown  precipitate  of  silver  oxide,  Ag2O,  is  formed. 
The  oxide  is  slightly  soluble  in  water,  and  the  solution  contains 
silver  hydroxide,  which  has  not  been  isolated;  it  is  a  relatively 
strong  base  and  yields  salts  which  are  not  hydrolyzed  by  water. 
The  oxide  dissolves  in  ammonia  and  forms  a  compound  of  the  com- 
position Ag(NH3)2OH,  which  is  converted  slowly,  but  more  rapidly 


COPPER,  SILVER,  AND  GOLD  591 

in  the  presence  of  caustic  alkalies,  into  a  black,  highly  explosive 
powder,  the  composition  of  which  is  unknown.  Silver  peroxide, 
Ag2C>2,  is  formed  as  the  result  of  the  action  of  ozone  on  silver. 

716.  The  Halides  of  Silver. — The  chloride,  bromide,  and  iodide 
of  silver  are  very  slightly  soluble  salts  and  are  obtained  as  curdy 
precipitates  when  the  respective  ions  are  added  to  a  solution  of 
a  silver  salt.     The  chloride  is  white,  the  bromide  pale  yellow,  and 
the  iodide  a  deeper  yellow;  the  solubilities  of  the  three  salts  are, 
respectively,  O.OslS,  O.C^l,  and  0.0e35  gram  in  100  grams  of  water 
at  18°. 

Silver  chloride  dissolves  readily  and  silver  bromide  less  readily 
in  a  solution  of  ammonia  and  form  salts  of  the  composition 
Ag(NH3)2Cl  and  Ag(NH3)2Br.  The  insoluble  salts  of  silver  dis- 
solve in  potassium  cyanide  and  form  a  soluble  complex  cyanide  of 
the  formula  KAg(CN)2.  Active  metals  like  zinc  displace  silver 
from  all  its  soluble  salts. 

Moist  silver  chloride  rapidly  changes  in  color  from  white  to 
violet  when  exposed  to  the  sunlight.  It  is  probable  that  the  color 
is  due  to  the  presence  in  the  salt  of  a  small  amount  of  colloidal 
silver,  which  is  produced  as  the  result  of  the  conversion  of  a 
part  of  the  chloride  under  the  influence  of  light  into  silver  and  free 
chlorine,  which  escapes.  The  halides  of  gold,  mercury,  lead,  and 
other  metals  are  also  reduced  in  the  presence  of  sunlight. 

717.  Silver  Nitrate. — This  salt  is  prepared  by  dissolving  silver 
in  nitric  acid.     It  forms  colorless  anhydrous  crystals,  which  melt 
at  208.6°.   Like  the  nitrates  of  the  alkali  metals,  it  is  more  stable 
when  heated  than  the  corresponding  salts  of  the  bivalent  metals. 
Advantage  is  taken  of  this  fact  in  purifying  silver  nitrate  pre- 
pared from  the  metal  which  contains  copper.     When  the  nitrates 
are  heated  the  copper  salt  is  converted  into  copper  oxide,  oxygen, 
and  oxides  of  nitrogen;   the  silver  salt  is  not  decomposed,  and  is 
extracted  from  the  mass  after  cooling. 

Silver  nitrate,  which  has  been  fused  and  cast  into  sticks,  is 
used  under  the  name  lunar  caustic  to  cauterize  sores.  The  salt 
is  also  used  in  some  forms  of  ink  to  mark  fabrics.  When  it  is  in 
contact  with  organic  material,  such  as  cotton  or  wool,  and  is 
heated  or  exposed  to  light,  it  is  reduced  and  metallic  silver  is 
deposited.  Inks  of  this  kind  should  not  be  used  in  marking 
clothing  that  is  to  be  washed  in  laundries,  because  the  silver  is 


592  INORGANIC  CHEMISTRY  FOR  COLLEGES 

converted  into  the  colorless  chloride  by  the  bleaching  solution 
commonly  used. 

718.  Test  for  Silver  Salts.— The  properties  of  the  halides  of 
silver  (716)  are  utilized  in  testing  for  the  metal.     It  is  precipitated 
as  chloride  in  the  first  group  in  qualitative  analysis  when  hydro- 
chloric acid  is  added  to  the  solution  to  be  analyzed.     The  chloride 
is  separated  from  lead  chloride  by  dissolving  the  latter  in  boiling 
water,  and  from  mercurous  chloride  by  treatment  with  ammonia, 
which  dissolves  silver  chloride.     It  is  obtained  from  the  solution 
made  in  this  way  by  adding  nitric  acid.     Black  silver  sulphide  is 
precipitated  when  solutions  of  silver  salts  are  treated  with  hydro- 
gen sulphide.     The  double  cyanides  containing  silver  are  not 
precipitated  in  the  presence  of  an  excess  of  an  alkali  cyanide. 
Silver  sulphide  is  reduced  by  nascent  hydrogen  to  silver. 

719.  Photography. — The  first  successful  photographic  process,  which  was 
discovered  by  Daguerre,  was  used  until  the  more  modern  methods  were 
developed.     Daguerrotypes  were  made  by  exposing  in  a  camera  a  plate  of 
silvered  copper  upon  the  surface  of  which  silver  iodide  had  been  deposited 
by  exposing  the  plate  to  the  vapor  of  iodine.     After  exposure,  the  plate  was 
set  over  a  vessel  containing  mercury,  which  was  gently  heated.     The  vapor 
of  the  metal  condensed  on  the  parts  of  the  plate  that  had  been  exposed  to 
light  and,  as  a  result,  the  picture  was  developed. 

In  modern  photography  the  material  of  the  plate  or  film  which  is  sensitive 
to  light,  is  prepared  by  making  a  jelly  from  gelatine,  in  which  is  precipitated 
silver  bromide  by  the  action  of  silver  nitrate  on  ammonium  bromide.  The 
jelly,  or  emulsion,  as  it  is  called,  is  first  solidified  and  converted  into  small 
pieces,  which  are  washed  thoroughly  with  cold  water  to  remove  the  soluble 
salts  present.  It  is  then  melted  and  heated  for  some  time  in  order  to  obtain 
the  colloidal  particles  of  the  silver  bromide  of  such  a  size  that  they  have 
the  desired  sensitiveness  to  light.  After  this  "ripening"  process,  the  melted 
emulsion  is  poured  on  plates  of  glass  or  celluloid  films  and  allowed  to  cool. 

A  great  deal  of  experim  entation  has  been  devoted  to  a  study  of  the  change 
that  takes  place  when  a  photographic  plate  is  exposed  to  light.  The  problem 
is  a  difficult  one  since  the  change  occurs  to  such  a  small  extent  that  its  effect 
cannot  be  seen  or  discovered  in  any  way  except  by  "development."  It  is 
believed  at  present  that  light  causes  the  decomposition  of  a  very  small  amount 
of  the  silver  bromide  into  silver  and  bromine,  which  is  absorbed  by  the  gel- 
atin. When  the  plate  is  treated  with  a  developer,  the  silver  bromide  which 
is  in  contact  with  the  trace  of  metallic  silver  set  free  by  the  light,  is  reduced 
to  the  metal,  which,  in  turn,  makes  it  possible  to  reduce  more  bromide. 
Silver  thus  serves  as  a  catalytic  agent  in  the  reduction  of  the  bromide,  and 
substances  are  used  as  developers  which  reduce  silver  bromide  very  slowly 
in  the  absence  of  the  metal.  In  the  early  days  of  photography  ferrous  salts 


COPPER,  SILVER,  AND  GOLD  593 

were  used  as  developers  because  they  reduced  silver  salts  to  the  metal  as 
the  result  of  their  oxidation  to  ferric  salts.  At  present  organic  compounds 
are  employed  almost  exclusively;  a  solution  of  pyrogallic  acid  in  an  alkali 
is  commonly  used  as  a  developer,  as  it  can  be  readily  oxidized. 

If  a  plate  has  been  overexposed,  it  is  reduced  very  rapidly  in  the  developer, 
and  an  unsatisfactory  negative  is  produced.  To  retard  development  and  thus 
allow  the  deposition  of  silver  in  such  a  way  that  normal  contrasts  between 
the  high  lights  and  shadows  are  produced,  a  trace  of  potassium  bromide  is 
added  to  the  developer;  this  reduces  the  solubility  of  silver  bromide  in  water 
(596)  and,  consequently,  the  rate  at  which  silver  is  deposited  from  the  salt. 

The  plate  is  next  "fixed"  by  placing  it  in  a  bath  containing  sodium  thio- 
sulphate,  commonly  called  "hypo,"  which  dissolves  from  the  plate  the  un- 
reduced silver  bromide.  The  reactions  which  take  place  are  as  follows: 

2AgBr+  Na2S2O3  =  Ag2S2O3  +  2NaBr 
Ag2S203  +  SNa&O,  =  2Na3Ag(S2O3)2 

Ordinary  alum  (688),  or  better,  chromium  alum,  K2S04,Cr2(SO4)3,  24H2O,  is  at 
times  added  to  the  fixing  bath,  and  is  of  especial  value  in  warm  weather,  as  it 
reacts  with  the  gelatin  and  hardens  it  so  that  it  holds  but  a  small  amount  of 
water,  and  can  be  dried  more  rapidly  and  with  less  danger  of  being  scratched. 

The  plate  is  next  washed  carefully  with  water  to  remove  all  of  the  thio- 
sulphates.  because  if  they  remain  on  the  plate  they  are  converted  on  stand- 
ing into  silver  sulphide,  which  forms  a  yellow  stain. 

A  positive  may  be  made  from  the  negative  by  printing  on  a  "gaslight" 
paper  which  has  been  coated  with  an  emulsion  resembling  in  composition 
that  used  on  the  plate;  it  usually  contains  some  silver  chloride,  and  is  much 
less  "rapid."  After  exposure  the  paper  is  developed  and  fixed  by  the  same 
method  used  with  plates.  Some  forms  of  paper  are  coated  with  silver  chloride, 
and  are  so  slow  in  action  that  they  can  be  handled  in  daylight.  The  print  is 
made  by  exposing  the  paper  under  the  negative  to  strong  sunlight  until  the 
picture  is  produced.  It  is  then  "toned"  by  placing  it  in  a  bath  containing 
a  gold  salt,  NaAuCl4;  the  silver  produced  as  the  result  of  the  action  of  light 
on  the  chloride,  passes  into  solution  and  gold  is  deposited  in  its  place.  The 
rate  at  which  the  latter  is  formed  determines  its  color,  which  may  vary  from 
brown  to  deep  purple.  The  desired  shade  is  obtained  by  adding  to  the  toning 
bath  substances,  such  as  borax,  which  affect  the  rate  at  which  the  metal  is 
deposited  and  the  size  of  the  colloidal  particles  formed.  Silver  prints  can 
also  be  toned  with  platinum.  The  so-called  platinum  papers  are  made  with- 
out the  use  of  silver  salts.  Ferric  oxalate,  Fe2(C2O4)3,  is  the  sensitive  salt 
used.  When  light  penetrates  the  negative  it  is  reduced  to  ferrous  oxalate, 
FeC2Od;  and  when  this  is  brought  into  contact  with  potassium  chloroplatinite, 
K2PtCl4,  the  latter  is  reduced  and  platinum  is  deposited.  Blue-print  paper  is 
coated  with  a  ferric  salt  as  its  light-sensitive  ingredient  and  with  potassium  fer- 
ricyanide,  which  reacts  with  the  ferrous  salt  formed  to  produce  a  blue  insoluble 
compound;  ferric-ammonium  citrate  is  commonly  used,  because  it  dissolves 
readilv  in  water  and  is  more  sensitive  than  most  other  ferric  salts. 


594  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  salts  of  silver  are  much  more  rapidly  affected  by  the  shorter  wave- 
lengths of  light,  and,  as  a  consequence,  a  print  from  a  negative  made  on  an 
ordinary  plate  does  not  show  the  gradation  in  brightness  of  the  original 
object.  Bright  shades  of  red  and  green  appear  dark,  and  deep  shades  of  blue 
appear  light.  This  effect  is,  in  part,  overcome  by  adding  to  the  emulsion 
used  on  the  plate,  dyes  that  absorb  the  longer  wave-lengths,  which  then 
become  active.  Plates  made  in  this  way  are  called  orthochromatic. 

COPPER 

720.  Copper  was  used  in  prehistoric  times  for  making  weapons 
and  tools,  and  later  was  alloyed  with  tin  to  form  bronze,  which 
was  the  most  important  metal  of  the  Greeks  and  Romans.     It 
was  replaced  for  these  purposes  by  iron  and  steel,  but  brass,  the 
alloy  of  the  metal  with  zinc,  proved  to  have  properties  which 
adapted  it  to  many  uses,  and  large  amounts  of  copper  are  used 
in  the  manufacture  of  the  alloy.     The  great  development  of  the 
electrical  industries  has  resulted  in  such  extensive  uses  of  the 
metal  that  it  now  ranks  next  to  iron  in  importance. 

In  the  stable  and  more  important  compounds  of  copper  the 
metal  has  the  valence  2,  but  it  also  forms  a  few  compounds  derived 
from  univalent  copper.  It  resembles  mercury  in  this  respect,  but 
the  series  of  compounds  derived  from  univalent  copper  is  less  com- 
plete than  that  of  the  mercury  compounds;  no  sulphate  or  nitrate 
of  cuprous  copper  has  been  isolated.  Copper  does  not  show  acid- 
forming  properties  as  zinc  does,  but  resembles  the  latter  in  uniting 
with  ammonia  to  form  complex  positive  ions,  for  example, 
Cu(NH3)4++.  The  salts  of  the  metal  are  slightly  hydroly zed  in 
solution  and,  like  zinc  salts,  yield  a  basic  carbonate  when  treated 
with  a  soluble  carbonate. 

721.  Occurrence  and  Metallurgy  of  Copper. — Copper  occurs  in 
a  large  number  of  minerals,  the  more  important  of  which  are 
chalcocite,  Cu2S,  chalcopyrite,  CuFeS2,  bornite,  CusFeSs,  cuprite, 
Cu2O,  azurite,  Cu3(OHJ2(CO3)2,  and  malachite,  Cu2(OH)2CO3. 
The  metal  also  occurs  in  the  free  condition  in  large  quantities  in 
the  Lake  Superior  mines  in  Michigan.     The  chief  ores  mined  in 
Montana  are  sulphides. 

On  account  of  the  increasing  demand  for  copper,  methods  have 
been  designed  to  use  as  ores,  materials  which  contain  as  low  as 
2  per  cent  of  copper  compounds  mixed  with  rock.  Such  ores  are 


COPPER,  SILVER,  AND  GOLD  595 

pulverized  and  "  concentrated,"  by  washing  them  with  water 
down  inclined  planes;  the  "gangue,"  being  lighter  than  the  ore, 
is  washed  away  faster,  and  a  material  containing  a  high  precentage 
of  the  latter  is  obtained.  This  method  has  been  largely  super- 
seded by  the  oil  flotation  process,  in  which  the  pulverized  ore  is 
agitated  with  water  in  the  presence  of  a  small  amount  of  oil.  The 
latter  adheres  to  the  ore  but  not  to  the  gangue,  and,  as  a  conse- 
quence, the  scum  which  rises  to  the  surface  contains  the  ore, 
which  can  be  readily  separated.  Oil  flotation  is  used  in  concen- 
trating ores  in  general,  and  has  made  it  possible  to  use  low- 
grade  material  which  could  not  be  economically  concentrated  by 
other  methods. 

The  methods  commonly  used  in  extracting  copper  from  its  ores 
have  been  described  briefly  in  the  chapter  in  which  metallurgy 
was  considered;  the  electrolytic  refining  of  copper  has  also  been 
discussed  (582). 

722.  Properties  of  Copper. — The  chief  physical  (table,  page 
443)  and  chemical  (549,  551)  properties  of  copper  as  well  as  its 
use  in  alloys  (542)  and  for  other  purposes  have  been  considered 
in  some  detail  in  previous  chapters.     The  metal  becomes  coated 
in  the  air  with  a  green  layer  of  a  basic  carbonate  and  is  attacked 
slowly  by  non-oxidizing  acids  if  oxygen  is  present.     A  deposit  of  a 
basic  chloride  is  formed  on  the  metal  when  it  is  left  in  contact  with 
sea- water. 

723.  The  Oxides  and  Hydroxides  of  Copper. — Cuprous  oxide, 
Cu2O,  occurs  in  nature  as  cuprite,  and  on  account  of  its  red  color 
is  called  ruby  copper.     It  is  made  by  reducing  an  alkaline  solution 
of  a  copper  salt  with  glucose.     The  precipitated,  hydrated  oxide, 
which  is  yellowish  red,  loses  most  of  its  water  quickly  and  yields 
the  anhydrous  red  oxide,  Cu2O,  when  gently  heated.     Like  silver 
oxide,  Ag2O,  it  dissolves  in  ammonia;    as  a  result  it  is  converted 
into  the  compound  of  the  formula  Cu(NHs)2OH,  which  is  colorless. 
The  solution  turns  blue  rapidly  if  air  is  present,  as  the  result  of 
oxidation    to    the    cupric    compound,    Cu(NH3)4(OH)2,    which 
is  deep  blue.     The  oxide  dissolves  in  concentrated  hydrochloric 
acid  and  a  compound  of  cuprous  chloride  and  the  acid  is  formed — 
H^CuCla.     Cuprous  salts  of  the  oxygen  acids  are  unknown,  and 
when  such  acids  are  added  to  cuprous  oxide,  a  cupric  salt  and  free 
copper  are  produced. 


596  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Cupric  oxide,  CuO,  is  formed  when  copper  is  heated  in  the  air, 
or  when  the  nitrate,  sulphate,  carbonate,  or  hydroxide  of  the  metal 
is  ignited.  It  is  an  active  oxidizing  agent  when  heated,  and  is 
used  as  such  in  the  analysis  of  compounds  containing  carbon 
and  hydrogen.  The  material  to  be  analyzed  is  heated  in  a 
stream  of  oxygen  in  a  hard  glass  tube,  and  the  vapors  produced  are 
passed  over  red-hot  copper  oxide,  which  insures  their  complete 
combustion.  The  water  formed  is  collected  in  a  tube  containing 
calcium  chloride,  and  the  carbon  dioxide,  in  a  bulb  containing  a 
solution  of  potassium  hydroxide. 

724.  Copper  hydroxide,   Cu(OH)2,   is  formed  as  a  light-blue 
precipitate  when  a  solution  of  a  base  other  than  ammonium 
hydroxide,  is  added  to  a  copper  salt.     If  the  solution  contains  a 
large  excess  of  alkali  and  it  is  heated,  the  hydroxide  loses  a  part  of 
its  hydrogen  and  oxygen  as  water  and  is  converted  into  a  black 
compound.     Cupric  hydroxide  dissolves  in  ammonia  and  forms  a 
deep-blue  solution,  which  contains  a  compound  of  the  formula 
Cu(NH3)4(OH)2.     Cotton,  paper,  and  other  forms  of  cellulose  dis- 
solve in  a  strong  solution  of  copper  hydroxide  in  ammonia.     The 
resulting  solution  is  used  in  making  one  form  of  artificial  silk  by 

'  forcing  it  through  capillary  tubes  into  a  dilute  solution  of  sulphuric 
acid,  which  precipitates  the  cellulose  in  the  form  of  a  fine  fiber. 
Other  kinds  of  artificial  silk  are  made  in  a  similar  way  by  using  a 
solution  prepared  by  the  action  of  sodium  hydroxide  and  carbon 
disulphide  on  cellulose  (viscose  silk),  or  a  solution  of  acetate  of 
cellulose.  The  products  made  in  this  way  are  not  silk,  which  is  a 
protein  and  contains  nitrogen,  but  cellulose  in  such  a  form  that 
is  has  a  high  luster.  Silk  dissolves  readily  in  a  warm  solution  of 
sodium  hydroxide,  but  artificial  silk  does  not. 

725.  Copper  hydroxide  dissolves  in  Rochelle  salt,  which  is 
sodium-potassium  tartrate,  NaK(C4H4O6).      The  blue  solution 
prepared  in  this  way  is  called  Fehling's  solution  and  is  used  in 
the  analysis  of  certain  sugars,  of  which  glucose  is  an  example. 
These  compounds  when  gently  heated  with  Fehling's  solution 
reduce  the  copper  to  hydrated  cuprous  oxide,  which  is  precipitated. 
The  reaction  is  utilized  in  testing  urine  for  glucose,  which  is  pro- 
duced in  the  body  and  is  excreted  in  the  case  of  the  disease 
known  as  diabetes, 


COPPER,  SILVER,  AND  GOLD  597 

726.  The  Halides  of  Copper. — Cupric  chloride  crystallizes  from 
water  as  a  blue  hydrate,  CuCl2,2H2O.  When  its  solution  is 
treated  with  ammonia,  the  green  basic  chloride,  CuCl2,3Cu(OH)2, 
first  precipitated,  dissolves  and  forms  a  deep-blue  solution  from 
which  a  compound  of  the  formula  Cu(NH3)4d2,H2O  can  be 
isolated. 

Cuprous  chloride,  CuCl,  is  prepared  by  heating  a  solution  of 
cupric  chloride  with  copper  in  the  presence  of  strong  hydro- 
chloric acid,  and  pouring  the  resulting  solution  into  water.  The 
cuprous  chloride  first  formed — CuCk  +  Cu  =  2CuCl — dissolves 
in  the  excess  of  the  acid  and  forms  a  soluble  compound  of  the 
formula  H^CuCls,  which  decomposes  when  the  solution  is  diluted: 
H^CuCls  =  CuCl  +  2HC1.  Cuprous  chloride  is  white,  insoluble  in 
water,  and  resembles  in  some  respects  silver  chloride  in  its  chemical 
properties.  It  is  soluble  in  ammonium  hydroxide  and  forms  a 
colorless  compound  of  the  formula  Cu(NH3)2Cl.  Moist  cuprous 
chloride  turns  violet  in  the  sunlight  as  the  result  of  reduction.  It 
is  oxidized  slowly  in  moist  air  to  a  green  basic  cupric  chloride. 
Ammoniacal  solutions  of  cuprous  chloride  or  other  cuprous  salts 
dissolve  carbon  monoxide,  and  unstable  addition-products  are 
formed,  the  composition  of  which  is  determined  by  the  pressure  of 
the  gas  in  the  solution.  The  reaction  is  used  in  the  quantitative 
analysis  of  flue  gases  and  in  the  purification  of  the  nitrogen  and 
hydrogen  used  in  the  synthetic  production  of  ammonia. 

Cupric  bromide,  CuBr2,  which  is  black  in  the  solid  condition, 
forms  concentrated  solutions  which  are  brown  but  change  on  dilu- 
tion to  the  blue  color  characteristic  of  solutions  of  copper  salts. 
The  change  in  color  is  probably  due  to  the  formation  of  copper  ions 
from  the  undissociated  salt. 

'Cupric  iodide  does  not  exist  at  ordinary  temperatures;  when 
a  soluble  iodide  is  added  to  a  solution  of  a  cupric  salt,  white  cuprous 
iodide  is  precipitated  and  iodine  is  set  free:  • 

2CuS04  +  4KI  =  2CuI  +  I2  +  2K2SO4 

The  reaction  is  used  in  the  quantitative  volumetric  analysis  of 
copper  compounds;  the  iodine  set  free  when  an  iodide  is  added  to 
the  solution  is  determined  by  titration  with  a  solution  of  sodium 
thiosulphate  (310). 


598  INORGANIC  CHEMISTRY  FOR  COLLEGES 

A  similar  reaction  takes  place  when  a  solution  of  potassium 
cyanide  is  warmed  with  a  copper  salt;  the  cupric  cyanide  first 
formed  decomposes  into  cuprous  cyanide,  CuCN,  and  cyanogen, 
(CN)2,  which  is  a  gas.  Cuprous  cyanide  dissolves  in  potassium 
cyanide  and  forms  the  complex  salt  KCu(CN)2,  which  is  very 
stable. 

727.  Copper  Sulphate. — This  salt,  which  is  called  blue  vitriol  or 
bluestone,  is  obtained  as  a  by-product  in  the  "  parting  "  of  gold 
and  silver,  and  also  by  allowing  dilute  sulphuric  acid  to  drip  on 
copper  scrap  in  the  presence  of  air.     In  another  process  the  scrap 
metal  is  heated  red  hot  in  a  furnace  and  sulphur  is  added  to  form 
copper  sulphide;  air  is  next  admitted  to  convert  the  latter  into 
a  mixture  of  oxide  and  sulphate,  which,  on  treatment  with  sul- 
phuric acid  and  water  and  subsequent  evaporation  yields  blue 
vitriol,    CuSO4,5H2O.     In   very   dry   air   the   crystals   effloresce 
and  change  to  a  white  powder,  which  has  the  composition  CuSCU,- 
H2O.     This  compound,  which  is  prepared  by  heating  the  hydrated 
salt  at  about  100°,  is  frequently  used  in  the  laboratory  for  drying 
liquids  which  react  with  calcium  chloride  or  potassium  hydroxide. 

Copper  sulphate  is  used  in  copper  plating,  as  a  mordant  in 
calico  printing  and  dyeing,  in  the  preparation  of  pigments  con- 
taining copper,  as  a  fungicide,  and  in  making  germicides  and  in- 
secticides. Bordeaux  mixture,  which  is  prepared  by  mixing  milk 
of  lime  and  a  solution  of  copper  sulphate,  is  used  to  spray  plants 
to  prevent  the  growth  of  fungi. 

728.  Other  Compounds  of  Copper. — Copper  nitrate  crystallizes 
from  cold  water  in  deliquescent  crystals,  Cu(N03)2,6H2O,  which 
are  converted  into  a  basic  salt  when  heated  at  100°.     The  basic 
copper    acetate   called  verdigris,  which  is  used  as  a  pigment,  is 
made  by  exposing  copper  to  acetic  acid  in  the  presence  of  air. 
The  formula   assigned  to  it  is  Cu(OH)2,2Cu(C2H3O2)2.     Paris 
green  is  prepared  by  heating  verdigris  with  a  solution  of  arsenious 
acid  containing  acetic  acid;    the  formula  assigned  to  the  light- 
green  product  formed  is  Cu(C2H3O2)2,Cu3(AsO3)2;   it  is  used  to 
exterminate  potato  bugs  and  insects.     Malachite  is  a  basic  car- 
bonate, CuCO3,Cu(OH)2,  found  in  nature,  which  is  used  as  a  green 
pigment.     A  compound  of  the  same  composition  is  formed  when 
sodium  carbonate  is  added  to  a  solution  of  a  copper  salt;  the 
reaction  resembles  that  which  takes  place  in  the  case  of  zinc  and  is 


COPPER,  SILVER,  AND  GOLD  599 

due  to  the  same  cause  (665).  Cuprous  sulphide,  Cu2S,  is  formed 
when  copper  is  heated  with  sulphur,  and  cupric  sulphide,  CuS, 
when  hydrogen  sulphide  is  passed  into  a  solution  of  a  copper  salt. 

729.  Tests  for  Copper  Salts. — Copper  is  precipitated  in  the 
second  group  in  qualitative  analysis  since  its  sulphide,  which  is 
black,  is  precipitated  by  hydrogen  sulphide  in  the  presence  of 
dilute  acids.     Copper  sulphide  dissolves  in  hot  dilute  nitric  acid. 
Soluble  cupric  salts  are  blue  in  solution  and  give  with  ammonium 
hydroxide  greenish  precipitates,  which  dissolve  in  an  excess  of 
the  reagent  and  produce  deep  blue  solutions  (595) .     Cupric  salts 
react  with  a  solution  of  potassium  ferrocyanide  and  give  cupric 
ferrocyanide,  which  is  a  red-brown  precipitate: 

2CuS04  +  K4Fe(CN)6  =  Cu2Fe(CN)6  +  2K2S04 

Copper  compounds  color  the  borax  bead  green  in  the  oxidizing 
flame  and  red  in  the  reducing  flame. 

GOLD 

Gold  is  univalent  in  some  of  its  compounds,  which  resemble 
the  analogous  derivatives  of  copper,  and  trivalent  in  others,  in 
which  it  functions  as  a  weak  acid-forming  element  and  resembles 
somewhat  aluminium  when  the  latter  acts  in  this  capacity. 

730.  Occurrence  and  Metallurgy  of  Gold. — Gold  occurs  chiefly 
in  the  metallic  condition  either  in  the  form  of  small  particles  mixed 
with  sand,  which  has  been  washed  into  the  beds  of  rivers,  or  dis- 
seminated through  veins  of  quartz  rock  which  have  not  been 
subjected  to  erosion.     It  occurs  in  this  form  in  the  Klondike,  the 
Transvaal  in  South  Africa,  and  California.     It  also  occurs  in  small 
quantities  in  the  sulphides  of  copper  and  lead,  and  is  recovered 
from  these  metals  in  their  purification  (582).     A  mineral  com- 
posed of  the  tellurides  of  gold  and  silver,  (Au,Ag)Te2,  is  mined  in 
Colorado;   the  metals  obtained  by  heating  it  to  drive  off  the  tel- 
lurium are  separated  by  treatment  with  strong  sulphuric  acid, 
which  dissolves  the  silver.     Since  the  acid  will  not  attack  the 
alloy  unless  it  contains  about  three-fourths  silver,  this  method  of 
"  parting  "  gold  and  silver  is  called  "  quartation."     The  silver 
sulphate  formed  is  dissolved  from  the  mixture  with  hot  water,  and 
the  metal  is  precipitated  by  means  of  metallic  copper.     The  gold 


600  INORGANIC  CHEMISTRY  FOR  COLLEGES 

left  as  a  powder  is  melted  in  a  cupel  with  potassium  nitrate  and 
borax,  if  base  metals  are  present. 

731.  Native  gold  is  extracted  from  the  mass  of  rock  with  which 
it  is  mixed  by  amalgamation  or  cyaniding.     The  rock  is  first 
broken  into  small  pieces  and  is  then  powdered  in  a  stamp-mill, 
which  consists  of  a  series  of  iron  mortars  supplied  with  cylindrical 
pestles  that  are  raised  and  lowered  by  machinery.     The  crushed 
ore,  mixed  with  water,  is  fed  into  the  mill  continuously  and  escapes 
through  a  fine  screen  containing  holes  about  0.5  mm.  in  diameter. 
It  next  passes  over  inclined  sheets  of  copper  coated  with  mer- 
cury,  from  which   the   amalgam   is   scraped   once   a   day.     The 
latter  is  then  pressed  in  chamois  skin  through  which  most  of  the 
mercury  passes;    the  hard  amalgam  left  is  distilled  in  a  retort. 
The  residue  is  refined  by  heating  with  niter  and  borax,   and 
"  parted." 

In  the  cyanide  process  the  finely  pulverized  ore  is  treated 
with  a  0.25  per  cent  solution  of  sodium  cyanide,  which  dissolves 
the  gold  in  the  presence  of  air: 

4Au  +  SNaCN  +  O2  +  2H20  =  4NaAu(CN)2  +  4NaOH 

The  solution  is  next  passed  through  boxes  containing  zinc  shavings, 
or  is  agitated  with  zinc  dust,  and  the  precipitate  formed  treated 
with  dilute  sulphuric  acid  to  dissolve  the  excess  of  zinc  and  the 
other  base  metals  present.  The  product  is  cupelled  and  parted. 

Gold  is  sometimes  obtained  by  chlorination.  In  this  process 
the  ore  is  treated  with  water  containing  chlorine  and  the  metal  is 
precipitated  by  ferrous  sulphate  from  the  chloride  which  is  formed: 

2AuCl3  +  6FeS04  =  2Au  +  2FeCl3  +  2Fe2(S04)3 
or  as  a  sulphide  by  hydrogen  sulphide : 

2AuCl3  +  3H2S  =  Au2S3  +  6HC1 

The  sulphur  is  removed  from  the  compound  by  heating  it. 

732.  Properties  of  Gold. — Pure  gold  is  bright  yellow,  but  the 
color  of  native  gold,  which  usually  contains  more  or  less  silver, 
varies  with  the  locality  from  which  it  is  obtained.     It  is  the  most 
malleable  and  ductile  of  the  metals,  and  can  be  made  into  sheets 
0.0002  mm.  thick;    in  this  form  it  is  used  as  gold-leaf  in  making 
signs  and  in  lettering  book  covers.  As  gold  is  quite  soft  it  is  alloyed, 


COPPER,  SILVER,  AND  GOLD  601 

usually  with  copper,  to  make  it  more  durable.  The  purity  of  gold 
is  designated  in  carats,  twenty-four  being  the  number  assigned 
to  the  pure  metal.  Jewelry  is  often  made  from  18-carat  gold, 
which  contains  ^f  of  the  metal,  but  on  account  of  the  softness  of 
this  alloy,  14-carat  gold  is  usually  preferred.  The  gold  coins  of 
Great  Britain  are  made  from  22-carat  gold,  the  other  metal  in  the 
alloy  being  copper.  Similar  alloys  are  used  in  the  coins  of  the 
United  States,  France,  and  Germany,  but  they  contain  90  per 
cent  of  gold,  which  is  equivalent  to  21.6  carats.  In  some  countries 
coins  are  made  from  an  alloy  of  gold  and  silver.  The  fineness  of 
gold  can  be  roughly  determined  by  rubbing  it  gently  with  a  fine- 
grained stone  and  testing  with  strong  nitric  acid  the  streak  of 
metal  formed;  the  result  is  compared  with  that  produced  with 
streaks  made  from  gold  of  known  composition. 

Gold  is  not  affected  by  the  substances  in  the  air  and  does  not 
decompose  hydrogen  sulphide.  It  does  not  dissolve  in  the  com- 
mon acids,  but  is  soluble  in  aqua  regia  as  the  result  of  the  forma- 
tion of  a  soluble  chloride. 

733.  Compounds  of  Gold. — Auric  chloride,  AuCls,  is  a  red 
crystalline  compound  formed  by  gently  heating  chlorauric  acid, 
HAuCU,  which  is  made  by  dissolving  gold  in  aqua  regia.  A  num- 
ber of  double  salts  are  derived  from  the  acid;  the  yellow  sodium 
salt,  NaAuCU,2Il2O,  is  used  in  photography  (719).  Auric  hy- 
droxide, Au(OH)3,  is  formed  as  a  yellowish  brown  precipitate 
when  a  solution  of  an  auric  compound  is  treated  with  sodium 
hydroxide;  it  dissolves  in  a  solution  of  sodium  hydroxide  and 
forms  a  salt  of  the  composition  NaAuO2,  which  is  derived  from 
metauric  acid,  Au(OH)3  —  B^O  =  HAuCb. 

Aurous  chloride,  AuCl,  is  formed  by  heating  auric  chloride 
at  about  180°;  it  is  colorless  and  is  converted  in  boiling-water  into 
auric  chloride  and  gold.  Aurous  iodide  is  precipitated  when  potas- 
sium iodide  is  added  to  a  solution  of  auric  chloride.  Aurous  oxide, 
Au2O,  is  a  violet  powder  which  reacts  with  hydrochloric  acid  to 
form  auric  chloride  and  gold.  It  will  be  seen  that  the  aurous 
compounds  resemble  closely  those  of  cuprous  copper.  Auric 
oxide,  Au2Oa,  is  a  brown  powder. 

Hydrogen  sulphide  precipitates  from  a  solution  of  auric  chloride 
a  brown  mixture  of  aurous  sulphide,  Au2$,  and  auric  sulphide, 
The  sulphides  dissolve  in  ammonium  sulphide  and  form 


602  INORGANIC  CHEMISTRY  FOR  COLLEGES 

thioaurites  and  thioaurates.  Potassium  cyanide  forms  two 
classes  of  complex  salts,  the  aurocyanides,  like  KAu(CN)2,  and 
the  auricyanides,  of  which  KAu(CN)4  is  an  example. 

734.  Analytical  Reactions  of  Gold. — The  behavior  of  soluble 
gold  compounds  with  hydrogen  sulphide  puts  the  metal  in  the 
analytical  group  with  arsenic,  antimony,  and  tin;  the  sulphides 
of  these  metals  are  precipitated  by  hydrogen  sulphide  in  the 
presence  of  strong  acids,  and  dissolve  in  yellow  ammonium  sul- 
phide. A  solution  of  gold  chloride  yields  a  characteristic  precip- 
itate of  the  metal,  called  purple  of  Cassius,  when  it  is  treated 
with  a  dilute  solution  of  stannous  chloride. 

EXERCISES 

1.  How  could  you    (a)  show  the  presence  of  Ag  in  an  aqueous  solution 
of  AgCl,   (6)  separate  AgCl  from  Agl,  and  (c)  prove  that  AgBr  contained  Br? 

2.  State  three  ways  of  distinguishing  AgNO3  from  KNO3  without  the 
use  of  chemicals. 

3.  How   could  you  separate  pure  Ag  from  a  mixture  of  AgNO3  and 
Pb(NO3)2? 

4.  Write  equations  for  reactions  by  which  Cu  can  be  obtained  from  CuSCX 
by  three  different  methods. 

5.  If  a  piece  of  silver  upon  which  has  been  deposited  a  thin  layer  of  Ag2>? 
is  brought  into  contact  with  a  piece  of  aluminium  under  a  dilute  solution  of 
sodium  hydroxide,  what  would  you  expect  to  happen?     For  what  purpose 
could  the  reaction  which  takes  place  be  used? 

6.  How  could  the  coating  that  forms  on  copper  in  the  air  be  removed? 

7.  How  could  you  obtain  in  two  different  ways  pure  copper  from  a  solution 
which  contains  CuSO4  and  A12(SO4)3? 

8.  How  could  you  distinguish  by  chemical  means  Cu  from  Cu2O? 

9.  Explain  what  you  think  would  happen  when  a  concentrated  solution 
of  hydrobromic  acid  is  added  to  an  aqueous  solution  of  CuBr2. 

10.  How  could  you  show  the  presence  of  Sb  and  Au  in  the  sulphides  pre- 
cipitated by  H2S,  which  dissolve  in  (NH4)2Sz? 

11.  When  a  sample  of  an  alloy  of  silver  and  copper  weighing  0.8572  gram 
was  dissolved  in  nitric  acid  and  the  silver  was  precipitated  as  chloride,  the 
latter  weighed  0.7167  gram.     Calculate  the  percentage  of  silver  in  the  alloy. 

12.  Calculate   what  weight   of     (a)  silver   chloride   and     (6) 
6H20  can  be  obtained  from  1  gram  of  silver  obtained  from  a  dime, 


CHAPTER  XLI 
IRON,  COBALT,  AND  NICKEL 

735.  Iron  was  used  in  the  earliest  times  of  which  we  have  any 
historical  records;  furnaces  for  the  metal  are  mentioned  by  Moses, 
and  the  art  of  making  weapons  from  iron  was  known  to  the  Egyp- 
tians and  Hindus.     The  introduction  in  1713  of  coke  in  place  of 
charcoal  as  the  agent  used  to  reduce  iron  ore,  the  utilization  of  a 
blast  of  air  in  1823  to  assist  in  obtaining  the  high  temperature 
required  in  the  furnace,  and  the  process  of  making  steel  invented 
by  Bessemer  in  1856  are,  perhaps,  the  three  most  important  devel- 
opments in  the  iron  industry  which  have  given  it  such  a  com- 
manding position  in  modern  life. 

736.  Occurrence  and  Metallurgy  of  Iron. — Iron  is  found  in  the 
metallic  condition  in  meteorites,  which  contain,  in  addition,  from 
3  to  8  per  cent  of  nickel  associated  in  certain  cases  with  smaller 
amounts  of  cobalt  and  copper.     The  metal  is  widely  distributed 
in  nature  in  the  form  of  silicates,  oxides,  and  the  sulphide,  FeS2, 
and  is  the  fourth  in  abundance  in  the  earth's  crust.     The  chief 
ores  of  the  metal  are  hematite,  Fe203,  limonite,  2Fe203,3H2O, 
magnetite,  Fe3O4,  and  siderite,  FeCOs. 

Iron  occurs  in  plants  and  animals  as  a  constituent  of  complex 
organic  compounds;  it  is  present  in  the  hemoglobin  of  the  blood, 
and  is  involved  in  this  condition  in  the  absorption  of  oxygen  in  the 
lungs. 

737.  Iron  is  obtained  by  reducing  its  ores  with  carbon.     The 
ores  contain,  in  addition  to  the  oxides  and  carbonate  of  the  metal, 
small    amounts    of    combined    sulphur,    phosphorus,    and    man- 
ganese, and  are  mixed  with  more  or  less  sand  and  clay.     In  order 
to  convert  the  latter  into  compounds  which  are  liquid  at  the  tem- 
perature at  which  reduction  is  carried  out  and  thus  make  it  pos- 
sible to  separate  them  from  the  metal  produced,  the  ore  is  mixed 
with  limestone;   this  furnishes  the  basic  material  to  flux  the  silica 

603 


604 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


and  alumina  by  converting  them  into  calcium  silicate  and  calcium 
aluminate.  The  reduction  of  the  ore  is  carried  out  in  a  blast 
furnace,  which  is  charged  with  a  mixture  of  the  ore,  calcium  car- 
bonate, and  coke.  Blast  furnaces  (Fig.  40)  are  constructed  of 
steel,  lined  with  brick,  and  are  from  80  to  100  feet  high,  and  from 
18  to  22  feet  in  diameter.  The  charge  is  carried  by  a  mechanical 
conveyor  to  the  top  of  the  furnace  and  introduced  into  it  through 
a  hopper  so  constructed  that  when  the  operation  is  carried  out 
there  is  little  loss  of  gas  from  the  furnace.  A  blast  of  highly 


FIG.  40. 

heated  air  enters  the  furnace  through  pipes,  called  the  tuyeres, 
and  rising  through  the  charge  converts  the  carbon  near  the  bottom 
into  carbon  dioxide.  The  latter  is  subsequently  reduced  by  the 
hot  carbon  above  it  to  carbon  monoxide,  which  passes  up  through 
the  furnace  and  reduces  the  oxide  of  iron. 

738.  There  are  four  zones  within  the  furnace  where  different  re- 
actions occur.  (Fig.  41 .)  In  the  upper  zone,  where  the  temperature 
is  from  100°  to  300°,  the  moisture  and  most  of  the  combined  water 
in  the  ore  are  removed;  in  the  next  at  300°  to  800°,  the  oxides  of 
iron  are  reduced  by  carbon  monoxide,  to  a  mixture  of  the  metal  and 
ferrous  oxide.  The  reaction  which  takes  place  between  an  oxide 


IRON,  COBALT,  AND  NICKEL 


605 


of  iron  and  carbon  monoxide,  for  example,  FeaCU  +  4CO  =  3Fe  + 
4CO2,  is  a  reversible  one  and  exothermic,  and,  as  a  consequence, 
the  equilibrium  is  displaced  unfavorably  for  the  reduction  with 
rise  in  temperature.  For  this  reason  a  large  excess  of  carbon 
monoxide  must  be  used  and  the  part  which  is  not  oxidized  by  the 
iron  escapes  with  the  other  gases  from  the  top  of  the  furnace.  In 

Coke  ^  i' me  stone 
Ore 


Oxygens 
Nitrogen-* 


JOSCO 

,3/eOi-COi     \   Blowing     HofKasf 
c  ~  Stains 


Coke.  Liquid 
Iron  and  Slqg_  _ 
~ 


Cement 

rfil/zer. 
Mineral  Woor. 


Wrouyfrffroo. 


FIG.  41. 


this  zone  the  limestone  and  siderite,  if  it  is  present,  are  converted 
into  oxides  and  carbon  dioxide. 

In  the  third  zone,  at  800°  to  1300°,  the  reduction  of  the  fer- 
rous oxide,  FeO,  is  completed  and  the  manganese,  silicon,  and 
phosphorus  are  converted  by  the  hot  carbon  into  the  elementary 
condition  and  unite  with  the  iron.  In  the  zone  of  fusion,  which 
comes  next,  the  temperature  is  from  1300°  to  1500°.  Here  the 


606  INORGANIC  CHEMISTRY  FOR  COLLEGES 

iron  melts  and  flowing  over  the  carbon  becomes  saturated  with  it, 
and  finally  drops  into  the  bottom  of  the  furnace.  In  this  zone  the 
lime  unites  with  the  siliceous  material  and  the  calcium  silicate 
and  calcium  aluminate  produced  melt  and  form  the  slag,  which 
floats  on  the  molten  iron. 

739.  Blast  furnaces  are  run  continuously  for  months;  the  slag 
is  tapped  off  through  a  hole  at  intervals  of  about  two  hours  and  the 
iron  every  four  to  six  hours.     The  composition  of  a  typical  slag 
is  30  to  35  per  cent  SiO2,  10  to  15  per  cent  A^Oa,  and  50  to  55  per 
cent  CaO.     Some  slags  are  ground  and  used  in  cement-making, 
and  others  not  adapted  to  this  purpose  are  used  as  ballast  on  rail- 
roads, for  road-building,  and  for  other  purposes.    The  molten  iron 
is  either  converted  directly  into  steel  or  run  into  molds  and  cast 
into  "  pigs." 

The  gas  which  issues  from  the  top  of  the  furnace  through  a  pipe 
called  the  "  downcomer  "  contains  over  20  per  cent  of  carbon 
monoxide,  and  possesses  a  high  heat-producing  value.  A  part 
of  it  is  used  to  heat  the  stoves  through  which  the  air-blast 
passes  before  being  admitted  to  the  furnace,  and  the  rest  is  freed 
from  dust  and  used  in  gas  engines  as  a  source  of  power. 

740.  Cast  Iron. — The  metal  obtained  from  the  blast  furnace 
is  known  as  pig  iron;   it  contains,  in  addition  to  the  metal,  from 
3  to  4  per  cent  carbon,  1  to  3  per  cent  silicon,  about  0.7  per  cent 
each  of  manganese  and  phosphorus,  and  from  0.02  to  0.05  per 
cent  sulphur. 

The  properties  which  are  characteristic  of  cast  iron  are  due 
largely  to  the  high  percentage  of  carbon  that  it  contains.  In  the 
liquid  metal  the  carbon  is  in  solution  in  the  form  of  a  carbide,  Fe3C, 
which  is  called  cementite.  The  properties  of  the  solid  iron  are 
determined  by  the  rate  at  which  the  metal  cools.  If  this  takes 
place  slowly,  as  in  a  sand  mold,  some  of  the  carbide  decomposes 
into  iron  and  carbon,  which  crystallizes  in  thin  black  scales  as 
graphite,  and  gray  cast  iron  is  obtained.  If  the  cooling  takes  place 
more  rapidly,  less  carbon  separates  and  white  cast  iron  is  formed. 
When  the  molten  iron  is  cast  in  a  mold  made  of  metal  the  part 
next  the  mold  is  suddenly  chilled  and  is  found  to  consist  of  a 
solid  solution  of  cementite  in  iron  without  any  graphite.  It  is 
called  chilled  cast  iron  and  is  very  hard  and  brittle.  Phosphorus 
increases  the  fluidity  of  the  molten  metal  and  aids  in  casting  it. 


IRON,  COBALT,  AND  NICKEL  607 

741.  Wrought  Iron. — The  properties  of  pure  iron  are  very  dif- 
ferent from  those  of  cast  iron;    it  softens  before  melting,  which 
takes  place  at  1530°,  whereas  cast  iron  melts  more  or  less  sharply 
at  temperatures  between  1150°  and  1250°,  which  depend  on  the 
amounts  of  the  other  elements  present.     The  purest  form  of  com- 
mercial iron  is  used  in  the  manufacture  of  piano  wire,  which  con- 
tains 99.7  per  cent  of  the  metal.     Wrought  iron,  which  is  so  called 
because  it  softens  on  heating  and  can  be  worked  when  hot,  is  made 
from  cast  iron  by  removing  most  of  the  other  elements  present 
in  it.     The  pigs  are  melted  on  the  hearth  of  a  reverberatory 
furnace,  which  is  covered  with  hematite  or  magnetite.     The  car- 
bon in  the  pig  iron  unites  with  the  oxygen  contained  in  the  oxides 
and  escapes  as  carbon  monoxide.     The  phosphorus  and  silicon 
are  also  oxidized  and  pass  into  the  slag  as  phosphates  and  silicates 
along  with  ferrous  sulphide  and  any  manganese  in  the  pig  iron. 
The  mixture  is  stirred  or  "  puddled,"  and  as  the  change  to  pure 
iron  progresses  it  becomes  stiffer  and  stiffer,  and  finally  is  worked 
into  balls  weighing  about  150  pounds,  which  consist  of  minute 
globules  of  iron  mixed  with  slag.     These  are  then  rolled  to  squeeze 
out  as  much  as  possible  of  the  slag  and  to  make  the  particles  of  iron 
coalesce.     This  process  leaves  about  2  per  cent  of  slag  in  the  metal 
which  contains  less  than  1  per  cent  of  other  impurities,  of  which 
about  0.2  is  carbon.     The  presence  of  slag  in  the  iron  causes  it  to 
have  a  fibrous  structure  which  is  said  to  increase  its  resistance  to 
breaking  and  to  assist  when  the  metal  is  welded.     Wrought  iron 
is  now  often  replaced  by  soft  steel,  which  resembles  the  former 
in  composition  closely  but  does  not  have  the  texture  which  results 
from  the  presence  of  slag. 

742.  Steel. — Cast  iron,  as  we  have  seen,  is  hard  and  brittle, 
whereas  wrought  iron  is  soft  and  can  be  welded;  the  first  contains 
a  relatively  high,  and  the  second  a  very  small  percentage  of  carbon: 
Steel  stands  between  these  two  in  properties  and  chemical  com- 
position, and  can  be  made  of  any  desired  degree  of  hardness  by 
regulating  its  content  of  carbon.     Soft  steels  that  can  be  welded 
contain  about  0.2  per  cent  of  carbon,  and  those  designed  for  struc- 
tural work  from  0.2  to  0.6  per  cent,  the  low  carbon  content  per- 
mitting the  rolling  of  the  metal  into  the  desired  form.     The  hard 
steels  required  for  tools  contain  from  0.9  to  1.5  per  cent  of  carbon. 
The  properties  of  iron  are  influenced  by  the  elements  other  than 


608 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


-  Flame 


carbon  in  pig  iron,  and  in  the  preparation  of  steel  from  the  latter 
most  of  the  sulphur,  phosphorus,  and  silicon  are  removed. 

In  converting  pig  iron  into  steel,  the  metal  is  largely  freed  from 
the  other  elements  present  by  oxidizing  them  in  the  presence  of 
substances  which  unite  with  the  oxides  formed  to  produce  a  slag. 
The  desired  amount  of  carbon  is  then  added  along  with  other  sub- 
stances which  are  used  to  remove  the  last  traces  of  oxygen  or 
oxides  present,  or  to  give  the  steel  the  special  properties  desired. 
A  number  of  ways  of  accomplishing  this  are  in  use,  the  more 
important  of  which  will  be  described  briefly. 

743.  The  Bessemer  Process. — In  this  process  as  used  in  the 
United  States  molten  pig  iron  is  poured  into  a  converter  (Fig.  42) 

which  is  constructed  of  steel,  lined 
with  sand  and  clay,  and  supported 
on  trunnions  so  that  it  can  be  rotated 
to  pour  out  the  finished  material.  Cold 
air  is  blown  through  the  molten  metal 
and  the  elements  present  in  the  iron 
are  converted  into  oxides;  the  silicon 
burns  first  and  unites  with  the  man- 
ganese oxide  formed  and  some  iron 
oxide  to  produce  silicates,  which  sepa- 
rate as  a  slag;  the  carbon  is  next 

oxidized  to  carbon  monoxide  which  burns  at  the  mouth  of  the 
converter.  Much  heat  is  generated  during  the  process  and 
the  temperature  is  kept  down  by  adding  solid  scrap  iron  or  by 
introducing  steam  with  the  blast.  In  about  fifteen  minutes, 
when  carbon  monoxide  no  longer  burns  at  the  mouth  of  the 
converter,  an  alloy  of  manganese,  iron,  and  carbon  is  added 
in  the  amount  necessary  to  produce  a  steel  of  the  required  com- 
position. 

Sulphur  and  phosphorus  are  not  removed  from  iron  in  the 
Bessemer  process,  because  at  the  temperature  of  the  converter 
iron  has  a  greater  affinity  for  these  elements  than  oxygen.  To 
make  steel  low  in  sulphur  and  phosphorus  it  is  necessary  to  use  pig 
iron  containing  not  more  than  0.1  per  cent  of  each  of  these  elements, 
or  to  line  the  converter  with  basic  substances  like  ignited  dolo- 
mite, which  will  take  up  sulphur  and  phosphorus  from  iron.  Basic 
linings  are  used  in  Europe,  but  in  the  United  States  steel  is  made 


FIG.  42. 


IRON,   COBALT,  AND  NICKEL 


609 


from  pig  iron  high  in  phosphorus  by  the  open-hearth  method 
which  is  described  below. 

744.  The  Open-hearth  Process. — In  this  process  pig  iron  is 
heated  on  a  hearth  the  lining  of  which  is  determined  by  the  amount 
of  the  elements  to  be  removed  from  the  iron.  If  the  metal  is  low 
in  phosphorus  and  sulphur,  the  lining  is  made  of  sand  or  other 
siliceous  material,  and  the  chemical  reactions  that  take  place 
resemble  those  in  the  Bessemer  converter;  this  is  known  as  the 
acid  open-hearth  process.  If  the  pig  iron  contains  more  of  these 
elements  than  can  be  present  in  the  steel,  the  lining  is  made  of 
calcined  magnesito  or  dolomite,  which  takes  up  the  phosphorus 
and  sulphur  from  the  iron.  The  basic  open-hearth  process  is 
commonly  used  in  the  United  States. 


FIG.  43. 

The  hearth  is  covered  with  a  roof  made  of  silica  brick  (Fig.  43) 
and  is  heated  by  gas  which  burns  directly  above  it.  The  charge, 
which  consists  of  pig  iron  and  enough  hematite  or  other  oxide  of 
iron  to  furnish  the  required  amount  of  oxygen  to  convert  the 
sulphur,  phosphorus,  and  silicon  into  oxides,  is  put  on  the  hearth 
and  heated  for  about  ten  hours.  In  order  to  obtain  as  high  a  tem- 
perature as  possible  the  gases  are  preheated  before  they  are 
brought  together  over  the  hearth.  This  is  accomplished  by  using 
what  is  called  a  regenerative  furnace  in  connection  with  the  hearth. 
The  hot  products  of  combustion  are  passed  through  two  chambers 
containing  " checker-work"  made  of  fire-brick,  until  they  become 
red  hot.  Gas  is  next  passed  through  one  of  the  chambers  and  air 
through  the  other  and  the  two  brought  together  over  the  hearth. 


610  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  products  of  combustion  are  passed  through  a  second  pair  of 
chambers  filled  with  checker-work.  By  shifting  at  intervals  the 
supply  of  air  and  gas  from  one  pair  to  the  other,  it  is  possible  to 
prevent  the  loss  of  the  large  amount  of  heat  which  under  ordinary 
conditions  passes  off  with  the  products  of  combustion. 

The  production  of  steel  in  this  way  is  slow  and  the  composition 
of  the  metal  can  be  tested  from  time  to  time  by  the  analysis  of  a 
sample  taken  from  the  hearth.  When  the  desired  changes  have 
taken  place,  the  material  is  drawn  off  into  ladles  to  which  are 
added  the  material  to  give  it  the  desired  composition.  Aluminium 
is  often  used  to  remove  the  dissolved  and  combined  oxygen,  and 
the  oxide  formed  passes  into  the  slag.  Titanium  is  also  used, 
because  it  unites  with  both  the  oxygen  and  the  dissolved  nitrogen. 
The  steel  made  with  these  deoxidizers  is  less  apt  to  break  under 
strain,  and  is  now  commonly  used  in  making  rails. 

745.  Crucible  Steel. — In  manufacturing  certain  kinds  of  steel 
the  process  is  carried  out  in  crucibles  made  of  a  mixture  of  graphite 
and  fire-clay  in  equal  amounts.     The  charge  consists  usually  of  a 
low-carbon  steel  or  wrought  iron  and  an  alloy  of  iron  and  manga- 
nese, either  f erromanganese  or  spiegel-eisen  (788) .     If  other  metals 
are  desired  in  the  steel,  such  as  nickel,  tungsten,  chromium,  vana- 
dium,  and  molybdenum,   they  are  also  added.     The   crucibles 
are  covered  and  heated  for  about  four  hours.     Steels  prepared  in 
this  way  are  used  in  making  tools,  drills,  dies,  files,  springs,  shafts, 
etc.     Large  quantities  of  steel  are  now  made  in  electric  furnaces. 

746.  Tempering  of  Steel. — The  physical  properties  of  steel 
are  determined  not  only  by  the  amount  of  carbon  and  other  ele- 
ments it  contains  but  by  the  heat  treatment  to  which  it  has  been 
subjected.     Steel  which  has  been  heated  to  redness  and  allowed 
to  cool  slowly  is  comparatively  soft;   if  it  is  plunged  into  water 
when  hot  it  becomes  very  hard.     Any  degree  of  hardness  between 
the  two  extremes  can  be  obtained  by  heating  the  chilled  metal  at 
different   temperatures  for   some   time.     This  process  is   called 
tempering,  and  the  temperatures  used  vary  with  the  purpose  to 
which  the  steel  is  to  be  put. 

747.  The  changes  which  take  place  in  tempering  steel    are 
complex  and  result  from  the  fact  that  pure  iron  can  exist  in 
three  distinct  modifications,  and  that  the  solubility  of  cementite, 
the  carbide  of  iron,  FesC,  which  is  present  in  steel,  varies  with  the 


IRON,  COBALT,  AND  NICKEL  611 

temperature.  Further,  the  changes  that  take  place  occur  slowly 
and  it  is  possible  by  more  or  less  rapid  cooling  to  pass  the  transition 
points  at  which  one  form  of  iron  changes  to  another;  in  the 
super-cooled  condition  of  the  metal  which  results,  a  variety  of 
forms  may  exist. 

Pure  iron  is  called  ferrite,  and  its  three  modifications  are  indi- 
cated by  prefixing  to  the  name  the  Greek  letters  alpha,,  beta,  and 
gamma.  a-Ferrite  is  the  form  of  iron  which  is  stable  below  760°. 
When  pure  iron  is  heated  a  change  occurs  at  this  temperature 
which  becomes  evident  by  the  heat  which  is  absorbed  without 
affecting  the  temperature;  it  is  a  true  transition  temperature 
where  a-ferrite  changes  into  /3-ferrite.  A  second  transition  occurs 
at  900°  when  7-ferrite  is  formed.  The  three  forms  differ  in 
properties;  the  first  can  be  magnetized,  the  other  two  not; 
7-ferrite  dissolves  up  to  1.7  per  cent  of  carbon  in  the  form  of 
cementite,  FeaC,  whereas  /3-ferrite  dissolves  much  less  and 
a-ferrite,  none  at  all.  If  steel  is  heated  above  900°  and  cooled 
rapidly,  there  is  not  time  for  the  cementite  to  separate,  and  the 
product  obtained  is  a  solid  solution  of  the  compound  in  7-ferrite, 
which  is  very  hard.  If  the  iron  contains  more  than  1.7  per  cent 
carbon,  the  excess  separates  as  graphite,  and  the  product  is  known 
as  pig-iron.  When  the  steel  cools  slowly  there  is  sufficient  time 
for  the  cementite  to  crystallize  out  and  for  the  7-ferrite  to  change 
to  a-ferrite.  Under  these  conditions  the  product  is  an  intimate 
mixture  of  iron  and  cementite,  and  is  comparatively  soft  if  it 
contains  less  than  0.8  per  cent  of  carbon. 

When  steel  is  heated  to  temper  it,  its  constituents  change 
slowly.  When  the  transformations  have  taken  place  to  such  a 
degree  that  the  metal  after  cooling  has  the  desired  hardness,  it  is 
more  or  less  rapidly  chilled  by  introducing  it  into  liquids  which 
take  up  the  heat;  if  the  cooling  is  to  be  very  rapid,  it  is  plunged 
into  water;  if  less  rapid,  into  oils  or  molten  salts. 

748.  Alloy  Steels. — By  alloying  steel  with  certain  metals, 
products  are  obtained  that  possess  properties  which  adapt  them  to 
special  uses.  The  effect  of  the  added  metal  is,  in  general,  to  lower 
the  transition  temperatures  at  which  the  different  forms  of  iron 
pass  one  into  the  other;  as  a  result,  7-ferrite  and  its  solid  solution 
of  cementite,  which  is  very  hard,  exist  at  lower  temperatures  than 
in  ordinary  steel.  For  example,  manganese  steel  which  contains 


612  INORGANIC  CHEMISTRY  FOR  COLLEGES 

from  12  to  14  per  cent  of  the  metal  and  about  1.5  per  cent  carbon, 
can  be  cast  and,  after  rapid  cooling,  is  very  hard.  It  is  used  in 
making  rock-crushing  machinery,  steam  shovels,  burglar  proof 
safes,  and  for  other  purposes  which  require  a  very  hard  steel. 
The  addition  of  a  very  small  amount  of  tungsten  or  molybdenum 
increases  the  hardness  of  the  metal. 

Nickel  steel  is  used  when  exceptional  toughness,  strength,  or 
hardness  is  desired  and  when  a  metal  is  required  to  resist  abrasion 
or  corrosion.  The  proportion  of  nickel  in  the  steel  varies  from  3.5 
to  42  per  cent.  Chromium  steels  are  also  hard  and  tough;  they 
are  used  for  files,  safes,  plows,  etc.  Very  small  amounts  of  vana- 
dium produce  about  the  same  effect  on  steel  as  that  produced  by 
nickel.  Steels  of  this  type  contain  approximately  0.15  per  cent 
vanadium,  0.25-0.5  per  cent  carbon,  0.5-1  per  cent  manganese  and 
1  per  cent  chromium  or  nickel;  they  have  a  high  elasticity,  are 
hard  and  tough,  weld  readily,  and  resist  shock  and  vibratory 
stresses.  For  this  reason  they  are  used  in  making  automobile 
parts.  High-speed  steels  do  not  lose  their  hardness  and  toughness 
at  red  heat,  and  are  used  in  making  tools  for  cutting  metals  at  such 
a  speed  that  the  friction  produced  heats  them  to  a  high  temperature. 
They  are  made  from  steel  to  which  15  to  20  per  cent  of  tungsten, 
3  to  5  per  cent  of  chromium,  and  a  small  quantity  of  vanadium  or 
cobalt  have  been  added.  An  alloy  of  iron  and  silicon  containing 
14  to  15  per  cent  of  the  latter,  which  is  sold  under  the  trade  name 
"duriron,"  is  used  in  making  castings  for  chemical  apparatus 
that  resists  to  a  high  degree  corrosion  by  acids. 

749.  The  Properties  of  Iron. — The  more  important  physical 
and  chemical  properties  of  iron  have  already  been  given.  When 
iron  burns  in  oxygen  or  is  heated  to  a  high  temperature  in  steam, 
it  is  converted  into  an  oxide  of  the  composition  FeaCU.  The 
reaction  with  steam  is  utilized  in  covering  iron  with  a  coating 
which  protects  it  against  rust  (552).  Passive  iron  and  the  action 
of  acids  on  the  pure  metal  have  been  discussed  (548,  547). 

Iron  forms  two  well-defined  series  of  compounds  in  which  it 
acts  as  a  base-forming  element;  the  ferrous  compounds,  in  which 
the  metal  has  the  valence  2,  resemble  the  compounds  of  mag- 
nesium and  are  hydrolyzed  to  but  a  slight  degree;  the  ferric  com- 
pounds contain  trivalent  iron  and  closely  resemble  the  analogous 
derivatives  of  aluminium.  Iron  acts  as  an  acid-forming  element 


IRON,  COBALT,  AND  NICKEL  613 

when  it  has  the  valence  6,  and  forms  salts,  called  ferrates,  which 
resemble  in  composition  the  sulphates. 

750.  Ferrous  Compounds. — When  iron  dissolves  in  dilute  acids 
it  is  converted  into  ferrous  salts.  Ferrous  chloride,  FeCl2,4H2O, 
forms  pale  blue  crystals,  which  turn  green  in  the  air  as  the  result  of 
oxidation  which  converts  them  into  a  basic  ferric  chloride.  In 
solution  the  salt  is  also  oxidized  by  air;  it  can  be  kept  in  the  ferrous 
condition  by  having  metallic  iron  and  a  small  amount  of  hydro- 
chloric acid  present.  In  order  to  obtain  ferrous  hydroxide  in  pure 
condition  the  solution  from  which  it  is  precipitated  by  an  alkali 
must  be  free  from  oxygen  and  ferric  salts.  This  is  best  accom- 
plished by  treating  with  a  boiling  solution  of  sodium  hydroxide, 
a  boiling  solution  of  ferrous  sulphate  which  contains  a  little  sul- 
phuric acid  and  is  in  contact  with  metallic  iron.  Under  these 
conditions  the  hydroxide  is  formed  as  a  white  precipitate.  It 
changes  rapidly  in  the  presence  of  air  to  a  green  color,  and  finally  is 
converted  into  ferric  hydroxide,  Fe(OH)3,  which  is  brown. 

Ferrous  carbonate,  FeCOs,  is  obtained  as  a  colorless  precipitate 
when  a  soluble  carbonate  is  added  to  a  ferrous  salt  in  the  absence 
of  oxygen.  It  is  soon  converted  in  the  presence  of  air  into  ferric 
hydroxide  as  the  result  of  oxidation  to  the  ferric  salt,  which  then 
hydrolyzes.  Ferrous  carbonate,  like  calcium  carbonate,  is  soluble 
in  water  containing  carbonic  acid  as  the  result  of  the  formation 
of  an  acid  salt,  Fe(HCO3)2-  It  is  present  in  this  form  in  hard 
waters  containing  iron,  which  deposit  ferric  hydroxide  as  the  result 
of  the  loss  of  carbon  dioxide  and  subsequent  oxidation  when  the 
water  comes  into  contact  with  the  air. 

Ferrous  sulphate,  FeSO4,7H2O,  which  is  also  called  green  vitriol 
or  copperas,  is  obtained  by  evaporating  a  solution  prepared  by 
dissolving  iron  in  dilute  sulphuric  acid.  The  salt  is  efflorescent 
and  becomes  brown  in  the  air  as  the  result  of  oxidation  to  a  basic 
ferric  sulphate,  Fe(OH)SO4.  The  double  salts  containing  ferrous 
iron  are  much  more  stable  in  the  air  than  the  simpler  compounds. 
Ferrous-ammonium  sulphate,  (NBU^SO^FeSO^GH^O,  crystal- 
lizes well  from  water  and  can  be  kept  in  a  pure  condition ;  it  is  used 
in  standardizing  solutions  for  use  in  quantitative  analysis.  Ferrous 
sulphate  is  used  as  a  mordant  in  dyeing,  as  a  disinfectant,  in  the 
purification  of  water  supplies,  in  the  manufacture  of  ink  and 
Prussian  blue,  and  in  metallurgy  for  precipitating  gold. 


614  INORGANIC  CHEMISTRY  FOR  COLLEGES 

Ferrous  sulphide,  FeS,  which  is  black,  is  made  by  heating  iron 
and  sulphur,  and  is  used  in  making  hydrogen  sulphide.  It  is 
formed  by  adding  ammonium  sulphide  to  a  solution  of  a  ferrous  salt. 

751.  Ferrous  cyanide,  Fe(CN)2,  is  precipitated  when  potassium 
cyanide  is  added  to  a  ferrous  salt.     It  dissolves  in  an  excess  of  the 
reagent  to  form  a  stable  complex cyanideof  theformulaK4Fe(CN)6,- 
3H2O,  which  is  called  potassium  ferrocyanide  or,  in  trade,  yellow 
prussiate  of  potash.     It  is  manufactured  by  heating  scrap  iron 
with  potassium  carbonate  and  animal  refuse  material  which  con- 
tains nitrogen  and  sulphur,  such  as  hair,  hoofs,  blood,  and  leather 
scrap.     Potassium  ferrocyanide  yields  K+    and    Fe(CN)e~" 
ions  when  dissolved  in  water.     The  fact  that  it  is  not  poisonous 
indicates  that  it  does  not  yield  any  potassium  cyanide.     Its  solu- 
tions give  none  of  the  reactions  characteristic  of  ferrous  salts. 
Potassium  ferrocyanide  precipitates  insoluble  ferrocyanides  from 
solutions  of  salts  of  most  metals.  Copper  ferrocyanide,  Cu2Fe(CN)  6, 
obtained  in  this  way  is  brown  and  serves  to  aid  in  the  identifica- 
tion of  copper  salts.     Ferrous  ferrocyanide  is  white,   but  soon 
turns    blue    as    the    result    of    oxidation.     Ferric     ferrocyanide, 
Fe4[Fe(CN)e]3,  is  a  deep-blue  precipitate  which  is  called  Prussian 
blue  and  is  used  as  a  pigment  and  in  making  laundry  blueing. 
Potassium  ferrocyanide  is  used  in  dyeing  and  calico  printing,  and  in 
making  potassium  cyanide  and  potassium  ferricyanide  (753). 

752.  Ferric    Compounds. — Ferric    chloride   in   the    anhydrous 
condition  is  prepared  by  heating  iron  in  a  stream  of  chlorine;   it 
sublimes  without   decomposition  and  forms  green  scales.     The 
salt  obtained  by  treating  ferrous  chloride  dissolved  in  water  with 
chlorine  and  evaporating  the  solution,  is  a  deliquescent  hydrate  of 
the  composition  FeCl3,6H2O.     The  yellow  color  of  a  solution  of 
ferric   chloride  is  largely  due  to   the   colloidal  ferric  hydroxide 
formed  as  the  result  of  hydrolysis.     The  color  can  be  intensified 
by  heating  the  solution  or  by  adding  to  it  a  small  amount  of  sodium 
hydroxide  to  neutralize  the  free  acid  present;   in  both  cases  the 
hydrolysis  is  increased.     Ferric  salts  are  converted  by  reducing 
agents  into  ferrous  salts.     Stannous  chloride  or  metallic  zinc  are 
commonly  used  for  this  purpose : 

2FeCl3  +  SnCl2  =  2FeCl2  +  SnCU 
2FeCl3  +  Zn  =  2FeCl2  +  ZnCl2 


IRON,   COBALT,  AND  NICKEL  615 

These  reactions  are  used  in  volumetric  analysis  in  the  quantitative 
determination  of  iron.  The  metal  is  reduced  to  the  ferrous  con- 
dition, and  the  volume  of  a  standardized  solution  of  an  oxidizing 
agent  required  to  convert  it  into  a  ferric  salt  is  determined. 

Ferric  hydroxide  is  formed  as  a  brown  precipitate  when  a  base 
is  added  to  a  solution  of  a  ferric  salt.  It  is  not  soluble  in  an  excess 
of  alkali.  It  is  readily  obtained  in  the  colloidal  condition  by 
suspending  in  a  stream  of  water  a  solution  of  ferric  chloride  con- 
tained in  a  bag  made  of  parchment,  which  is  an  animal  mem- 
brane. The  molecules  of  hydrochloric  acid  which  result  from  the 
partial  hydrolysis  of  the  chloride  are  of  such  size  that  they  can  pass 
through  the  minute  pores  in  the  parchment,  but  the  larger  mole- 
cules of  the  ferric  hydroxide  cannot.  As  the  acid  diffuses  through 
the  membrane  the  equilibrium  between  it  and  the  hydroxide  is 
destroyed— FeCl3  +  3H20  ;=±  Fe(OH)3  +  3HC1— and  finally  all  of 
the  salt  is  hydrolyzed.  Prepared  in  this  way  ferric  hydroxide 
remains  in  colloidal  solution — that  is,  it  is  apparently  dissolved 
in  the  water,  but  it  shows  none  of  the  properties  of  an  electrolyte ; 
it  has  no  effect  on  the  freezing-point  and  does  not  conduct  an 
electric  current.  The  solution,  known  as  dialyzed  iron,  has  a  red 
color  and  is  used  as  an  antidote  in  arsenical  poisoning. 

Ferric  hydroxide  is  precipitated  when  a  carbonate  is  added  to  a 
solution  of  a  ferric  salt,  the  reaction  being  analogous  to  that  which 
occurs  in  the  case  of  aluminium  salts.  When  solutions  containing 
the  salts  of  bivalent  -and  trivalent  metals  are  shaken  with  solid 
barium  carbonate,  the  hydroxides  of  the  trivalent  metals  only  are 
precipitated.  The  reaction  serves,  therefore,  as  a  means  of  sep- 
arating these  two  kinds  of  metals  and  is  used  in  qualitative 
analysis. 

Ferric  oxide,  Fe20s,  is  extensively  used  as  a  red  pigment.  It 
is  prepared  by  treating  with  sodium  carbonate  the  liquors  obtained 
in  cleaning  the  surface  of  iron,  and  igniting  the  precipitate  formed. 
In  this  form  it  is  sold  under  the  name  rouge.  If  sulphuric  acid 
has  been  used  for  "  pickling,"  the  solution  is  first  neutralized 
with  lime;  the  precipitate,  which  contains  ferric  hydroxide  mixed 
with  calcium  sulphate,  yields  on  ignition  a  light  red  pigment 
called  Venetian  red. 

The  oxide  of  iron  of  the  formula  FeaCU,  which  occurs  in  nature, 
is  called  lodestone  or  magnetic  oxide  on  account  of  the  fact  that  it 


616  INORGANIC  CHEMISTRY  FOR  COLLEGES 

will  attract  iron.  It  is  formed  by  heating  iron  in  oxygen  or  super- 
heated steam,  and  is  a  compound  of  ferrous  and  ferric  oxides. 

Ferric  sulphate  is  made  by  oxidizing  with  nitric  acid  a  solution 
of  ferrous  sulphate  containing  sulphuric  acid.  It  does  not 
crystallize  well  and  is  very  soluble  in  water.  Like  the  sulphates 
of  the  other  trivalent  metals  it  forms  alums. 

753.  The  addition  of  a  cyanide  to  a  ferric  salt  causes  the  pre- 
cipitation of  ferric  cyanide,  which  dissolves  in  an  excess  of  the 
reagent  and  forms  potassium  ferricyanide,  K3Fe(CN)e  [Fe(CN)3, 
3KCN],  which,  on  account  of  its  color,  is  called  red  prussiate  of 
potash.     It  is  manufactured  by  treating  a  solution  of  potassium 
ferrocyanide  with  chlorine: 

2K4Fe(CN)6  +  C12  =  2K3Fe(CN)6  +  2KC1 

When  a  solution  of  potassium  ferricyanide  is  added  to  one  of  a 
ferrous  salt,  a  precipitate,  called  Turnbull's  blue,  is  formed: 

3FeSO4  +  2K3Fe(CN)6  =  Fe3[Fe(CN)6]2  +  3K2SO4 

With  ferric  salts  the  ferricyanide  gives  no  precipitate,  but  a  brown 
solution  is  formed.  The  use  of  the  ferricyanide  in  photography 
has  been  mentioned  (719) . 

754.  Iron  as  an  Acid-forming  Element. — Ferric  hydroxide  does 
not  dissolve  in  solutions  of  bases,  but  when  it  is  fused  with  caustic 
alkalies  it  reacts  with  the  base  and  salts,  called  ferrites,  are  pro- 
duced.   These  salts,  of  which  sodium  ferrite,  NaFe02,  is  an  exam- 
ple are  derived  from  the  compound  formed  from  ferric  hydroxide 
by  the  loss  of  water,  Fe(OH)3  —  H20  =  HFeO2;   they  are  com- 
pletely hydrolyzed  by  water,  which  produces  the  base  and  ferric 
hydroxide:    NaFeO2  +  2H2O  =  NaOH  +  Fe(OH)3.     This   reac- 
tion has  been  utilized  in  a  commercial  process   (Loewig's)   for 
the  manufacture  of  caustic  soda  and  yields  a  very  pure  product. 
Iron  oxide  is  heated  at  a  high  temperature  with  sodium  carbonate 
and  the  product,  after  being  washed  with  cold  water,  is  decom- 
posed by  hot  water.     The  iron  hydroxide  is  recovered  and  used 
again. 

If  a  solution  of  potassium  hydroxide  in  which  precipitated 
ferric  hydroxide  is  suspended,  is  treated  with  chlorine,  the  iron  is 
oxidized  and  converted  into  potassium  ferrate,  K2FeO4,  which  is 
obtained  in  the  form  of  red  crystals  on  evaporating  the  solution. 


IRON,  COBALT,  AND  NICKEL  617 

The  salt  is  hydrolyzed  by  water  in  the  absence  of  an  excess  of 
alkali,  and  the  ferric  acid  formed  decomposes  into  ferric  hydroxide 
and  oxygen.  In  the  ferrates,  which  resemble  the  sulphates  in 
composition,  iron  has  the  valence  6  as  has  sulphur  in  sulphates. 
The  metal  can  be  oxidized  to  this  valence  only  in  the  presence  of 
strong  alkalies  with  which  the  acid  formed  can  unite  to  form  a  salt. 

755.  The  Corrosion  of  Iron. — The  changes  that  take  place 
when  iron  rusts  have  been  discussed  briefly  (550),  but  their  fuller 
consideration  was  delayed  until  the  properties  of  the  oxides 
and  hydroxides  of  the  metal  and  the  hydrolysis  of  its  salts  had  been 
studied,  because  all  these  are  involved  in  the  process.  On  account 
of  the  industrial  importance  of  the  corrosion  of  iron,  the  changes 
involved  in  it  have  been  carefully  studied  in  order  to  arrive  at  an 
explanation  of  the  process,  upon  which  methods  to  prevent  cor- 
rosion can  be  based.  A  number  of  views  as  to  the  way  in  which 
iron  rusts  have  been  advanced,  but  the  one  based  upon  electro- 
chemical considerations  is  now  generally  accepted.  According 
to  an  older  view  carbon  dioxide  was  necessary  in  ordinary  rusting; 
it  united  with  the  water  present  to  form  carbonic  acid,  which 
reacted  with  iron  and  formed  ferrous  carbonate  and  hydrogen. 
The  salt  was  next  hydrolyzed  by  water,  and  the  ferrous  hydroxide 
produced  was  oxidized  by  the  air  to  ferric  hydroxide;  the  carbon 
dioxide  liberated  served  to  react  with  more  iron,  and  the  process 
continued. 

In  the  more  recent  explanation,  the  presence  of  carbon  dioxide 
is  not  considered  necessary,  although  it  increases  the  rate  of  cor- 
rosion in  the  manner  just  indicated.  In  the  presence  of  pure  water 
iron  passes  into  solution  to  a  slight  degree  as  ferrous  hydroxide, 
and  hydrogen  is  liberated.  In  the  presence  of  the  air  the  dis- 
solved hydroxide  is  oxidized  to  ferric  hydroxide,  which  separates 
from  solution  on  account  of  its  very  small  solubility.  Pure  iron 
rusts  very  slowly,  but  commercial  iron  and  steel,  which  contain  a 
number  of  other  substances,  corrode  rapidly.  According  to  the 
electrochemical  explanation  of  the  process,  the  other  substances 
present  in  the  iron  form  with  it  an  electric  couple,  of  which  the 
metal  is  the  more  positive  constituent.  Local  action  is  set  up  in 
the  way  already  described  in  the  case  of  metallic  couples  (567), 
and  the  iron  is  rapidly  converted  into  ferrous  hydroxide,  which  is 
oxidized  to  the  hydrated  ferric  oxide.  Iron  does  not  rust  evenly, 


618  INORGANIC  CHEMISTRY  FOR  COLLEGES 

but  at  certain  places  on  its  surface,  and  pitting  occurs  where 
there  is  a  segregation  of  the  impurities  in  the  metal.  The  presence 
of  "  mill  scale/'  which  is  the  oxide  formed  on  the  surface  of  the 
hot  metal  when  it  is  worked,  markedly  increases  the  rate  at  which 
iron  rusts.  The  scale  is  brittle  and  when  it  cracks  and  leaves  a 
surface  of  iron  exposed,  a  couple  is  set  up  and  the  metal  is  cor- 
roded. Rust  itself  forms  a  couple  with  pure  iron  and  when  it  is 
once  formed  the  corrosion  proceeds  more  rapidly. 

It  has  been  shown  that  the  rate  of  rusting  during  the  second 
year  is  about  twice  as  fast  as  during  the  first  year.  Unused  rail- 
road rails  rust  more  rapidly  than  those  from  which  the  rust  is 
removed  by  the  jarring  produced  by  the  passage  of  cars  over 
them.  When  iron  is  strained  in  any  way  it  assumes  a  different 
potential  from  that  of  the  unstrained  metal,  and,  as  a  result, 
corrosion  is  greater  in  the  neighborhood  of  punched  holes  than 
around  holes  drilled  in  the  metal.  Corrosion  takes  place  where 
the  metal  has  been  scratched  by  a  file  or  struck  with  a  heavy  tool. 

Since  oxygen  is  involved  in  the  corrosion  of  iron,  the  metal 
rusts  very  slowly  when  it  is  immersed  in  deep  water,  which  con- 
tains but  a  little  of  the  gas.  Rain  water,  which  is  saturated  with 
oxygen  and  carbon  dioxide,  affects  iron  very  rapidly. 

756.  Tests  for  Iron  Salts. — Solutions  of  ferrous  salts  give  a 
deep-blue  precipitate  with  potassium  ferricyanide  (753)  and  a 
black  precipitate  of  ferrous  sulphide,  which  is  soluble  in  acids,  on 
the  addition  of  a  solution  of  ammonium  sulphide.  Ferric  salts 
give  a  deep-blue  precipitate  with  potassium  ferrocyanide  (751) 
and  an  intense  red  coloration  with  ammonium  thiocyanate, 
NH4SCN.  When  sodium  acetate  is  added  to  a  solution  of  a  ferric 
salt,  a  red  solution  of  ferric  acetate  is  formed,  which  deposits,  when 
heated,  an  insoluble  basic  acetate  that  results  from  the  hydrolysis 
of  the  salt.  The  reaction  is  utilized  in  the  quantitative  separation 
of  iron  from  manganese  in  the  analysis  of  steel.  The  acetates  of 
the  trivalent  metal  are  converted  into  insoluble  basic  acetates 
when  heated  in  water  solution,  whereas  those  of  the  bivalent 
metals,  which  are  stronger  base-forming  elements,  are  not  decom- 
posed in  this  way.  Iron  salts  color  the  borax  bead  green  in  the 
reducing  flame  (ferrous  borate)  and  yellow  in  the  oxidizing  flame 
(ferric  borate). 


IRON,  COBALT,  AND  NICKEL  619 


COBALT 

757.  Cobalt  occurs  as  smaltite,  CoAs2,  and  cobaltite,  CoAsS, 
and  is  associated  with  nickel  in  its  ores.     The  metal  has  been  used 
recently  in  high-speed  tool  steels  (748),  and  it  has  been  suggested 
as  a  substitute  for  nickel  in  plating  other  metals.     Cobalt  forms 
two  series  of  compounds  in  which  it  has  the  valence  2  and  3, 
respectively.     The  cobaltous  compounds  are  but  slightly  hydro- 
lyzed  and  resemble  those  derived  from  bivalent  iron.      Cobaltous 
chloride,  CoCl2,6H20,  crystallizes  from  water  in  red  prisms,  which 
become   blue   when   dehydrated.     Solutions   of    caustic   alkalies 
precipitate  from  solutions  of  cobaltous  compounds  blue  basic  salts 
that  are  converted  by  boiling  into  cobaltous  hydroxide,  Co(OH)2, 
which  is  pink,  when  first  formed,  but  changes  to  brown  as  the  result 
of  oxidation  to  cobaltic  hydroxide,  Co  (OH)  3.     Cobaltous  hydrox- 
ide dissolves  in  a  solution  of  ammonia  to  form  the  compound 
Co(NH3)4(OH)2,  which  is  rapidly  oxidized  to  the  cobaltic  com- 
pound, Co(NHs)6(OH)3,  which  is  blue.     Cobaltous  sulphate  has 
the  composition  CoSO4,7H2O,  and  cobaltous  nitrate,  Co(NO3)2,- 
6H2O;    both  salts  are  red.     The  sulphide,  CoS,  is  black  and  is 
precipitated  by  ammonium  sulphide. 

758.  When  a  solution  of  sodium  hypochlorite  is  added  to  a 
solution  of  a  cobaltous  salt,  cobaltic  hydroxide  is  formed  as  a  black 
precipitate.     The  hydroxide  dissolves  in  cold  hydrochloric  acid 
with  the  formation  of  cobaltic  chloride,  which  decomposes  on 
warming    into    cobaltous    chloride    and    chlorine.     The    cobaltic 
salts  are  all  highly  hydrolyzed  in  water,  but  are  more  stable  in  the 
form  of  double  salts.     Of  these  the  complex  cyanides  are  the  most 
important.     Potassium   cobalticyanide,    K3Co(CN)6,    is   prepared 
by  the  action  of  chlorine  on  potassium  cobaltocyanide,  K4Co(CN)e. 
The  compounds  resemble  closely  the  analogous  salts  containing 
iron. 

A  potash-cobalt  glass  is  used  under  the  name  smalt  as  a  pig- 
ment and  in  making  blue  glazes  for  china.  It  is  prepared  by 
fusing  together  sand,  potassium  carbonate,  and  cobaltic  oxide, 
Co2Os.  Cobalt  blue  is  an  excellent  pigment  made  by  igniting  a 
mixture  of  alumina,  and  basic  cobalt  phosphate  previously  pre- 
pared by  the  action  of  sodium  phosphate  on  a  solution  of  cobalt 
nitrate. 


620  INORGANIC  CHEMISTRY  FOR  COLLEGES 

759.  Tests  for  Cobalt  Salts. — When  cobaltous  salts  are  treated 
in  solution  with  acetic  acid  and  potassium  nitrite,  they  are  first 
oxidized  to  the  cobaltic  condition  by  the  nitrous  acid  set  free, 
and  are  then  converted  into  a  complex  nitrite  of  the  formula 
KsCo(NO2)6,  which  is  formed  as  a  yellow  precipitate.    The  reaction 
is  characteristic  of  cobalt  and  is  used  as  a  test  for  the  element  in 
qualitative  analysis.     Cobalt  colors  the  borax  bead  blue  in  both 
the  oxidizing  and  reducing  flame. 

NICKEL 

760.  Nickel  occurs  in  the  metallic  condition  in  meteorites. 
The  chief  source  of  the  world's  supply  of  the  metal  is  the  Sudbury 
district  in  Ontario,  where  an  iron  sulphide  occurs  which  contains 
about  2  per  cent  each  of  nickel  and  copper.    A  hydrated  silicate 
of  magnesium,  nickel,  and  iron  is  used  as  an  ore  in  New  Caledonia. 
The  ores  are  smelted  and  the  nickel  and  copper  separated.     In  the 
Monde  process  the  reduction  is  carried  out  in  a  Bessemer  converter, 
and  the  nickel  is  then  converted  into  nickel  carbonyl,  Ni(CO)4, 
by  heating  the  metal  at  about  50°  in  a  stream  of  carbon  mon- 
oxide.    The  product,  which  is  a  gas,  is  passed  through  cylinders 
heated  to  200°  in  which  nickel  is  deposited  as  the  result  of  the 
decomposition  of  the  carbonyl  at  this  temperature. 

Nickel  is  separated  electrolytically  from  copper,  or  the  alloy 
which  is  the  result  of  the  reduction  of  the  ore  is  used  as  such 
under  the  name  monel  metal,  which  is  stronger  than  ordinary 
steel  and  resists  the  action  of  acids. 

Nickel  is  used  in  alloys  (542),  for  plating  (583),  and  in  the 
production  of  certain  steels;  about  60  per  cent  of  the  world's 
production  is  used  for  the  latter  purpose  (748).  Invar,  a  steel 
which  contains  30  per  cent  nickel,  has  a  heat-coefficient  of  expan- 
sion of  practically  zero  and  is  used  in  making  pendulums,  etc. 
Nickel  coins  contain  1  part  of  the  metal  alloyed  with  3  parts  copper. 

761.  Nickel  forms  two  oxides,   NiO  and  Ni20a.      Nickelous 
hydroxide,  Ni(OH)2,  is  converted  by  acids  into  nickelous  salts, 
which  are  stable  in  the  air  and  are  not  oxidized  to  nickelic  com- 
pounds.    Nickelic  hydroxide,  Ni(OH)s,  has  even  less  basic  pro- 
perties than  cobaltic  hydroxide;  it  is  formed  as  a  black  precipitate 
when  a  hypochlorite  is  added  to  a  solution  of  a  nickel  salt. 


IRON,  COBALT,  AND  NICKEL  621 

The  chloride,  NiCl2,6H20,  and  sulphate  of  nickel,  NiS04,7H2O, 
are  green  salts  which  crystallize  from  water.  The  double  sulphate 
of  the  formula  (NH4)2Ni(SO4)2,6H2O  is  used  in  nickel  plating  (748). 
Nickelous  hydroxide,  which  is  green,  is  obtained  by  precipitation; 
it  dissolves  in  ammonia  and  forms  a  colorless  solution  of  a  com- 
pound of  the  formula  Ni(NH3)4(OH)2. 

762.  Test  for  Nickel  Salts. — When  potassium  cyanide  is  added 
to  a  solution  of  a  nickel  salt  the  green  nickelous  cyanide,  Ni(CN)2, 
first  precipitated  dissolves  in  an  excess  and  forms  a  complex 
cyanide  of  the  formula  K2Ni(CN)4,H2O,  which  does  not  resemble 
in  composition  and  properties  the  double  cyanides  containing 
iron  and  cobalt  in  the  bivalent  condition.  It  is  less  stable  and 
yields  black  nickelic  hydroxide  when  treated  with  hypochlorites. 
The  reaction  serves  to  separate  nickel  from  cobalt  since  potas- 
sium cobalticyanide  is  not  affected  by  hypochlorites.  Nickel 
differs  from  cobalt  in  that  it  does  not  form  an  insoluble  double 
nitrite. 

EXERCISES 

1.  How  could  you  distinguish  from  one  another  the  following  ores  of  iron: 
(a)  hematite,    (6)  limonite,    (c)  magnetite,  and   (d)  siderite? 

2.  What  would  happen  if  oxygen  -were  passed  over  metallic  iron  con- 
tained in  a  silica  tube  that  was  heated  red  hot? 

3.  A  large  excess  of  CO  is  required  in  reducing  the  oxides  of  iron  in  a  blast 
furnace.     What  conclusion  can  you  draw  in  regard  to  the  reaction  FeO  -f-  CO 
=  Fe  +  CO2?     Is  the  conclusion  in  accord  with  the  facts? 

4.  What  chemical  compounds  would  you  expect  to  be  present  in  a  blast- 
furnace slag? 

5.  Name  some  articles  made  of    (a)  cast  iron,     (6)  wrought  iron,     (c) 
high  carbon  steel,  and  state  a  reason  in  each  case. 

6.  Write  equations  for  the  reactions  which  occur  in   (a)  the  acid  Bessemer 
process,    (6)  the  basic  open  hearth  process. 

7.  (a)  Why  cannot  FeCl3  be  made  by  heating  FeCl3,6H2O?     (6)  What 
would  be  obtained  if  the  latter  salt  were  heated  to  a  high  temperature? 

8.  Write  an  equation  for  the  reaction  which  takes  place  when  an  aqueous 
solution  of  FeSO4  containing  H2SO4  is  heated  with  HNO3. 

9.  How  could  you  test   (a)  a  ferric  salt  for  the  presence  of  a  ferrous  salt, 
and   (6)  a  ferrous  salt  for  the  presence  of  a  ferric  salt?     Write  equations  for 
the  reactions  used. 

10.  Could  Fe2O3  which  contains  A12O3  be  used  in  Loewig's  process  for 
making  NaOH?     Give  a  reason  for  your  answer. 

11.  Write  an  equation  for  the  reaction  which  occurs  when  K2FeO<  is 
dissolved  in  water. 


622  INORGANIC  CHEMISTRY  FOR  COLLEGES 

12.  Is  it  a  good  practice  to  set  iron  fence  posts  in  stone  by  placing  the 
post  in  a  hole  in  the  stone  and  filling  the  cavity  with  lead?     What  improve- 
ment could  you  suggest? 

13.  What  would  be  the  result  if  an  iron  pipe  laid  in  the  ground  remained 
in  contact  for  a  long  time  with  a  wire  carrying  an  electric  current?     Explain. 

14.  (a)  Out  of  what  form  of  iron  should  anchors  be  made?     Give  reasons. 
If  an  iron  anchor  is  used  to  hold  a  mooring  in  place  and  remains  under  water 
a  long  time  what  would  you  expect  if   (6)  the  water  is  shallow  and  the  bottom 
is  rocky,    (c)  the  water  is  shallow  and  the  bottom  is  covered  with  growing 
plants,    (d)  the  water  is  deep  and  the  bottom  covered  with  soft  mud? 

15.  Write  equations  for  the  reactions  which  take  place  when  a  solution 
of  ferric  chloride  is  treated  with  sodium  acetate  and  then  heated.     Assume 
that  the  basic  acetate  formed  has  the  formula  Fe(OH)2(C.,H3O2). 

16.  Show  by  a  consideration  of  the  properties  of  the  compounds  of  Ni 
and  Co  which  of  the  elements  is  the  more  closely  related  to  Fe.     Are  the 
atomic  weights  of  the  elements  in  accord  with  your  conclusion? 


CHAPTER  XLII 
THE  PLATINUM  METALS 

763.  The  metals  in  the  eighth  group  in  the  periodic  classification 
other  than  iron,  cobalt,  and  nickel  are  classed  together  under  the 
name  platinum  metals,  because  they  are  found  associated  with 
platinum  and  resemble  the  latter  in  physical  properties.     They 
occur  as  alloys  which  frequently  contain  more  or  less  gold,  and  are 
obtained  as  nuggets  and  small  particles  by  washing  the  alluvial 
sands  deposited  by  certain  rivers.     The  world's  supply  is  obtained 
almost   exclusively   from   the   Ural   Mountains,   although   small 
amounts  of  platinum  are  separated  from  the  gold  found  in  Australia, 
Brazil,  and  California,  and  from  the  copper-nickel-iron  deposits 
of  Sudbury,  Ontario. 

764.  Platinum. — In  separating  platinum,  which  is  the  most 
important  metal  of  the  group,   the   naturally  occurring   alloys 
are  treated  with  aqua  regia,  and  to  the  solution  of  the  chloride 
formed  is  then  added  ammonium  chloride,  which  precipitates  a 
double  chloride  of  the  formula  (NH^PtCle-     This  salt  on  being 
strongly  heated  decomposes  and  leaves  a  residue  of  platinum. 

Platinum  is  a  soft,  very  heavy  metal  which  has  the  specific 
gravity  21.4  and  melts  at  1755°.  Its  resemblance  to  silver  in 
appearance  led  the  Spaniards,  who  discovered  it  in  South  America, 
to  give  it  the  name  platinum,  which  is  the  diminutive  of  the 
Spanish  word  for  silver,  plata.  The  fact  that  the  metal  is  com- 
paratively inactive  chemically  has  adapted  it  to  many  uses.  It 
has  recently  been  used  extensively  in  making  jewelry,  and  it  is 
claimed  that  it  accentuates  the  beauty  of  diamonds  when  they  are 
set  in  it,  but  its  popularity  can,  no  doubt,  be  traced,  in  part,  to  the 
fact  that  at  present  it  is  much  more  costly  than  gold. 

The  most  important  uses  of  the  metal  are  in  chemical  industries 
and  in  laboratories.  In  the  manufacture  of  sulphuric  acid  it 
finds  extensive  application  as  a  contact  agent  in  the  preparation 
of  sulphur  trioxide,  and  in  the  construction  of  evaporating  pans 

623 


624  INORGANIC  CHEMISTRY  FOR  COLLEGES 

used  in  making  the  concentrated  acid.  It  is  the  most  important 
catalyst  used  in  the. oxidation  of  ammonia  to  nitric  oxide,  which 
is  converted  into  nitric  acid.  The  metal  is  used  in  the  laboratory 
in  the  form  of  resistance  wire,  electrodes,  crucibles,  dishes,  etc. 
It  was  formerly  used  in  large  quantities  in  the  construction  of 
electric-light  bulbs  and  other  electrical  appliances  in  which  it 
is  necessary  to  pass  a  metallic  conductor  through  glass.  The 
variation  of  the  coefficient  of  expansion  of  platinum  with  temper- 
ature is  about  the  same  as  that  of  glass,  and,  as  a  consequence, 
the  latter  does  not  crack  where  the  metal  has  been  sealed  into 
it.  Platinum  has  been  largely  replaced  for  this  purpose  by  wire 
made  of  nickel-steel  coated  with  copper. 

765.  Platinum  forms  compounds  in  which  it  shows  the  valencies 
2  and  4.     When  it  has  the  higher  valence  it  has  acid-forming 
properties.     For  this  reason  the  metal  is  attacked  when  in  contact 
with  fused  alkalies,  especially  if  oxidizing  agents  are  present.     It 
is  not  affected  by  fused  alkaline  carbonates  and,  consequently, 
crucibles  of  paltinum  are  used  in  the  laboratory  in  preparing  sub- 
stances for  analysis  which  are  decomposed  by  molten  carbonates. 
Red-hot  platinum  allows  hydrogen  to  pass  through  it  freely,  and 
unites  with  arsenic,  lead,  phosphorus,  etc.     For  this  reason  sub- 
stances that  contain  elements  which  are  reducible  at  red  heat  by 
hydrogen  should  not  be  ignited  in  a  platinum  crucible.     Platinum 
also  unites  slowly  with  carbon  at  high  temperatures  and  becomes 
so  brittle  that  it  is  apt  to  crack. 

766.  When  platinum  is  dissolved  in  aqua  regia  it   is  converted 
into   chloroplatinic   acid,   H^PtClejOH^O,   which   forms    reddish-- 
brown deliquescent  crystals.     The  acid  yields  an  ammonium  salt, 
(NH^PtCle,  and  a  potassium  salt,  K^PtCle;    the  latter,  which 
is  difficultly  soluble  in  water,  is  used  in  the  quantitative  determi- 
nation  of   potassium.     When   chloroplatinic   acid   is   heated   at 
about  240°  it  is  converted  into  platinous  chloride,  PtCl2,  which 
is  a  green  compound,  insoluble  in  water,  that  dissolves  in  hydro- 
chloric acid  and  forms  chloroplatinous  acid;   the  potassium  salt, 
K^PtCU  is  used  in  photography  (719) .    The  double  platinocyanides 
of  potassium  and  barium,  K2Pt(CN)4,3H20  and  BaPt(CN)4,4H2O, 
are  employed  in  making  screens  for  X-ray  work;    paper  coated 
with  either  of  these  salts  glows  brightly  when  exposed  to  the  rays 
on  account  of  the  fact  that  the  compounds  fluoresce,  that  is,  they 


THE  PLATINUM  METALS  625 

convert  the  invisible  radiation  of  short  wave-lengths  into  those  of 
longer  wave-lengths  that  produce  the  sensation  of  light. 

767.  Bases  precipitate  as  a  black  powder  from  solutions  of 
platinous  chloride,  platinous  hydroxide,  Pt(OH)2,  which  dissolves 
in  acids  but  not  in  bases.     Platinic  hydroxide,  Pt(OH)4,  which  is 
yellow,  is  formed  in  the  same  way  from  platinic  compounds;    it  is 
soluble  in  both  acids  and  bases,  and  with  the  latter  yields  plati- 
nates.     Both  sulphides  of  the  metal,  PtS  and  PtS2,  dissolve  in 
ammonium  polysulphide  and  form  ammonium  thioplatinates. 

768.  The   Other  Metals   of  the   Platinum   Group. — A   large 
number  of  the  compounds  of  the  metals  associated  with  platinum 
have  been  described,  but  cannot  be  considered  here.     Ruthenium 
and  osmium  form  oxides  of  the  formulas  RuCU  and  OsC>4,  and  show, 
therefore,  the  valence  8,  in  accordance  with  their  position  in  the 
periodic   classification.     The   elements  have   acid-forming   prop- 
erties and  yield  salts  of  the  composition  K2RuO4  and  K2OsO4. 
The  former,  potassium  ruthenate,  is  converted  by  large  amounts 
of  water  into  an  oxide  of  the  metal  and  potassium  perruthenate, 
KRuO4,  in  which  ruthenium  has  the  valence  7.     The  tetroxide  of 
osmium,  OsCU,  is  commonly  called  "  osmic  acid/'  although  it  does 
not  form  salts;   it  is  used  in  histology  for  staining  and  hardening 
tissues.     Osmium  melts  at  about  2700°  and  was  formerly  used 
in  making  filaments  for  electric  lamps  but  has  been  replaced  for 
this  purpose  by  tungsten.     Wires  made  of  an  alloy  of  platinum 
and  rhodium  are  used  in  one  form  of  thermocouples.     Rhodium 
is  also  used  in  making  tips  for  gold  pens.     An  alloy  of  iridium  and 
osmium  is  also  used  for  the  latter  purpose.     Iridium  is  used  in  wires 
for  thermocouples  and  is  usually  present  in  commercial  platinum. 
The  latter  when  pure  is  a  relatively  soft  metal  and  a  small  per- 
centage of  iridium  makes  it  harder  and,  therefore,  more  durable. 

Palladium  is  characterized  by  its  ability  to  absorb  large  quan- 
tities of  hydrogen  and  is,  therefore,  a  valuable  catalytic  agent  in 
the  union  of  the  gas  with  other  substances  which  react  with  hydro- 
gen. One  volume  of  the  metal  in  the  form  of  foil  will  absorb  500 
volumes  of  the  gas,  and  when  precipitated  as  a  fine  powder,  about 
800  volumes. 

The  most  characteristic  compounds  of  the  elements  of  this 
group  are  the  double  chlorides  which  they  form  with  the  chlorides 
of  the  alkali  metals. 


CHAPTER  XLIII 
CHROMIUM,  MOLYBDENUM,  TUNGSTEN,  AND  URANIUM 

769.  The  consideration  of  the  members  of  the  first  family  of 
the  sixth  group  in  the  periodic  classification  of  the  elements  has 
been  delayed  to  this  point  on  account  of  the  fact  that  the  most 
important  member  of  the  family,  chromium,  forms  compounds  in 
which  it  plays  the  parts  of  a  bivalent  and  a  trivalent  metal  and 
others  in  which  it  acts  as  an  acid-forming  element.  The  chemistry 
of  the  compounds  of  chromium  can  now  be  studied  in  the  light  of 
the  facts  learned  in  regard  to  the  behavior  of  bivalent  and  trivalent 
metals. 

The  chromates,  of  which  potassium  chromate,  E^CrCU,  is  an 
example,  resemble  closely  in  properties  the  sulphates.  The 
similarity  in  their  physical  structure  is  evident  from  the  fact 
that  potassium  chromate  is  isomorphous  with  potassium 'sulphate, 
and  sodium  chromate,  Na2CrO4,10H2O,  with  Glauber's  salt,  which 
is  the  hydrate  of  sodium  sulphate  of  analogous  composition. 

It  will  be  recalled  that  sulphuric  acid  and  sulphur  trioxide 
are  oxidizing  agents  on  account  of  the  fact  that  the  valence  of 
sulphur  can  change  from  6  to  4  or  even  less.  In  the  case  of  chromic 
acid  and  its  anhydride,  CrOs,  the  reduction  in  valence  takes  place 
so  readily  that  these  compounds  find  many  applications  in 
chemistry  as  active  oxidizing  agents.  When  acting  in  this  way 
chromium  changes  from  the  valence  6,  in  which  it  is  an  acid- 
forming  element,  to  the  valence  3,  in  which  it  is  metallic  in  chemical 
properties  and  resembles  aluminium. 

The  other  elements  of  the  family  form  trioxides,  which  are  acid 
anhydrides,  and  their  most  important  compounds  are  salts  which 
resemble  the  sulphates  in  composition. 

CHROMIUM 

770.  The  most  important  ore  of  chromium  is  chromite,  or 
chrome-iron  ore,  FeO,Cr2O3,  which  resembles  magnetite, 

626 


CHROMIUM,  MOLYBDENUM,  TUNGSTEN,  AND  URANIUM     627 

FeO,Fe2O3,  in  composition;  it  is  considered  to  be  a  ferrous  salt  of 
an  acid  derived  from  chromic  hydroxide  by  the  loss  of  water — 
Cr(OH)s  —  H2O  =  HCrO2 — and  its  chemical  name  is  ferrous 
chromite.  The  element  also  occurs  as  crocoisite,  which  is  lead 
chromate,  PbCrC>4,  and  in  traces  in  emeralds  and  other  precious 
stones,  to  which  it  gives  their  characteristic  colors.  Chromium 
may  be  obtained  by  reducing  the  trioxide,  Cr2O3,  with  aluminium, 
but  the  large  amount  of  the  metal  used  in  making  steel  (748)  is 
obtained  by  the  reduction  of  chromite  by  carbon  in  the  electric 
furnace;  the  alloy  of  chromium  and  iron  obtained  in  this  way  is 
used  directly. 

771.  Properties   of  Chromium. — The   element  is   crystalline, 
white,  very  hard,  melts  at  1520°,  boils  at  2200°,  and  has  the 
specific  gravity  6.9.     It  is  not  affected  by  the  air,  but  burns  in 
oxygen  and  forms  the  trioxide,  Cr2O3.     It  dissolves  slowly  in 
hydrochloric  acid  and  chromous  chloride,  CrCl2,  is  formed;  its 
behavior  with  nitric  acid  resembles  that  of  iron,  and  like  the 
latter  it  is  rendered  passive  by  the  concentrated  acid  (548) .     The 
chief  use  of  metallic  chromium  is  in  the  preparation  of  alloys  (748). 

772.  Chromous  Compounds. — Chromous  chloride  can  be  made 
as  stated  above,  or  by  heating  chromic  chloride  in  hydrogen.     It  is 
also  formed  by  the  action  of  metallic  zinc  on  a  solution  of  chromic 
chloride:    2CrCl3  +  Zn  =  2CrCl2  +  ZnCl2.     The  anhydrous  salt 
is  white  and  its  solution  in  water  is  blue.     It  is  very  unstable  and 
in  the  presence  of  hydrochloric  acid  is  rapidly  oxidized  to  chromic 
chloride  by  the  oxygen  of  the  air:  4CrCl2  +  4HC1  +  O2  =  4CrCl3 
-f  2H2O.     Chromous  hydroxide,  formed  by  the  action  of  bases  on 
chromous  salts,  is  a  yellow  precipitate,  which  is  such  a  powerful 
reducing  agent^  that  it  reacts  slowly  with  water  and  liberates 
hydrogen:  2Cr(OH)2  +  2H2O  =  2Cr(OH)3  +  H2.    Chromous  sul- 
phate, CrSO4,7H2O,  resembles  closely  in  chemical  and  physical 
properties  ferrous  sulphate,  FeSO4,7H20,  but  it  is  a  much  more 
active  reducing  agent. 

773.  Chromic  Compounds. — The  large  number  of  salts  derived 
from  trivalent  chromium  resemble  closely  those  of  aluminium  and 
ferric  iron.     Chromic  hydroxide,  Cr(OH)s,  is  formed  by  the  action 
of  bases  on  chromic  salts;  it  is  pale  blue  in  color  and  reacts  with 
acids  to  form  chromic  salts,  and  with  a  large  excess  of  a  caustic 
alkali  to  form  a  chromite;  potassium  chromite,  KCrO2,  formed  in 


628  INORGANIC  CHEMISTRY  FOR  COLLEGES 

this  way  is  derived  from  an  acid  produced  as  the  result  of  the  loss  of 
water  from  the  hydroxide  (754).  The  chromites  are  highly  hydro- 
lyzed,  and  when  heated  with  water  are  converted  into  chromium 
hydroxide,  which  is  precipitated:  KCrO2  +  2H20  =  Cr(OH)3  + 
KOH.  Chromic  hydroxide  is  formed  when  chromium  salts  are 
used  as  mordants  (688)  and  in  tanning  (688),  and  is  the  compound 
which  brings  about  the  desired  results  in  these  important  processes. 
Chromic  chloride)  CrCl3,6H2O,  is  obtained  in  the  form  of  bluish- 
gray  crystals  by  the  evaporation  at  room  temperature  of  a  solution 
of  chromic  hydroxide  in  hydrochloric  acid.  When  dissolved  in 
water  the  salt  gives  a  violet  solution,  which  changes  to  green  when 
heated  to  boiling. 

774.  From  this  solution  can  be  isolated  a  second  form  of 
chromic  chloride,  which  is  green  but  has  the  same  composition  as 
that  of  the  blue  salt.     Silver  nitrate  precipitates  all  the  chlorine 
from  the  blue  salt,  but  only  one-third  of  that  contained  in  the  green 
chloride.     This  difference  in  behavior  is  explained  on  the  hypoth- 
esis that  the  blue  chloride  gives  the  ions  [Cr(H2O)e]  +  +  "    and 
3C1~  and  the  green  chloride  the  ions  [CrCl2(H2O)4]+  and  Cl~. 

775.  Chromic  oxide,  C^Os,  is  used  as  a  pigment  under  the 
name  chrome  green.      It  is  prepared  by  igniting  at  red  heat  pre- 
cipitated chromic  hydroxide. 

Chrome  alum,  K2SO4,Cr2 (864)3, 24H2O,  forms  deep  violet 
crystals,  which  behave  when  heated  with  water  in  a  way  similar 
to  that  mentioned  above  in  the  case  of  chromic  chloride.  The 
salt  gives  a  green  solution  on  boiling,  which  deposits  a  gummy 
mass  on  evaporation  and  from  which  barium  chloride  does  not 
precipitate  a  sulphate.  On  standing  for  some  time  in  cold  water, 
the  green  salt  changes  and  the  violet  form  is  deposited  as  crystals. 
Chrome  alum  is  used  as  a  source  of  chromic  hydroxide  in  tanning 
and  as  a  mordant. 

776.  The  Chromates. — The  salts  of  chromic  acid,  which  are 
used  extensively  as  oxidizing  agents  and  as  pigments,  are  manu- 
factured from  potassium  chromate;  the  latter  is  made  by  roasting 
chromite  FeO,Cr2Os,  with  potassium  carbonate  and  lime  in  the 
presence  of  air.     Under  these  conditions  the  iron  in  the  mineral  is 
oxidized  to  ferric  oxide,  Fe2O3,  and  the  chromium  to  chromic 
anhydride,  CrOs,  which  reacts  with  the  carbonate  to  form  potas- 
sium chromate,  K2CrO4.     The  lime  which  is  added  to  keep  the 


CHROMIUM,  MOLYBDENUM,  TUNGSTEN,  AND  URANIUM     629 

mixture  porous  and  thus  allow  the  air  to  act  upon  it,  is  also  con- 
verted, in  part,  into  calcium  chromate.  The  product  is  treated 
with  water  to  dissolve  out  the  chromates,  and  enough  potassium 
sulphate  is  added  to  precipitate  the  calcium  as  sulphate. 

Potassium  chromate  is  an  anhydrous  salt  and  forms  pale 
yellow  crystals  which  are  very  soluble  in  water.  Lead  chromate, 
PbCrO4,  has  a  brilliant  yellow  color  and  is  the  basis  of  the  "  chrome 
yellows  "  used  as  pigments,  which  are  made  by  mixing  it  with 
other  insoluble  salts  such  as  the  sulphates  of  lead,  barium,  or  cal- 
cium; it  turns  black  in  the  presence  of  hydrogen  sulphide.  The 
solubilities  of  the  chromates  are  very  similar  to  those  of  the  cor- 
responding sulphates. 

777.  Chromic  Anhydride  and  the  Dichromates. — When  con- 
centrated sulphuric  acid  is  added  to  a  strong  solution  of  potassium 
chromate,  chromic  anhydride  is  precipitated  in  the  form  of  dark 
red  needles:  K2CrO4  +  H2SO4  =  K2SO4  +  H2O  +  CrO3.  The 
chromic  acid,  which  is  probably  first  formed,  decomposes  into  its 
anhydride  and  water.  The  compound  dissolves  readily  in  water, 
but  differs  from  sulphuric  anhydride,  SO3,  in  that  but  a  small 
proportion  of  it  unites  with  water  to  form  an  acid.  It  does  unite, 
however,  with  neutral  chromates  to  form  stable  compounds  which 
are  called  dichromates:  K2CrO4  +  Cr03  =  K2Cr2O7. 

When  the  correct  amount  of  sulphuric  acid  is  added  to  a  solu- 
tion of  a  chromate,  a  dichromate  is  formed;  one-half  of  the 
salt  ?s  converted  into  potassium  sulphate  and  chromic  anhy- 
dride, and  the  rest  unites  with  the  latter  to  form  the  dichromate: 

K2CrO4  +  H2SO4  =  K2SO4  +  [CrO3]  +  H2O 
K2CrO4  +  [CrO3]  =  K2Cr2O7 


2K2CrO4  +  H2SO4  =  K2SO4  +  K2Cr2O7  +  H2O 

When  potassium  hydroxide  is  added  to   a  solution  of   potassium 
dichromate  it  is  converted  into  the  chromate: 

K2Cr2O7  =  K2CrO4  +  [CrO3] 
2KOH  +  [CrO3]  =  K2CrO4  +  H2O 


K2Cr207  +  2KOH  =  2K2Cr04  +  H20 


630  INORGANIC  CHEMISTRY  FOR  COLLEGES 

The  change  is  evident  on  account  of  the  fact  that  solutions  of 
dichromates  are  red  and  those  of  chromates  are  light  yellow. 

Potassium  dichromate,  which  is  orange-red  in  color,  can  be 
readily  purified  by  crystallization  from  water  on  account  of  its 
relatively  small  solubility  at  room  temperature  (8  parts  in  100  of 
water  at  10°).  It  was  for  this  reason  that  it  was  formerly  pre- 
pared in  large  quantities  by  the  method  described  above  (776) 
and  was  used  in  making  other  compounds  of  chromium.  It  has 
recently  been  replaced  largely  by  the  sodium  salt,  Na2Cr2O7,2H2O, 
which  forms  red  crystals  and  is  very  soluble  in  water  (109  parts  in 
100  at  15°). 

Potassium  dichromate  ionizes  largely  to  produce  K+  and 
Cr2<37~~  ions  but  the  solution  also  contains  free  CrOa  and  CrO4~" 
ions  formed  as  the  result  of  the  slight  decomposition  of  the  dichro- 
mate into  the  chromate  and  chromic  anhydride.  It  is  for  this 
reason  that  if  potassium  dichromate  is  added  to  a  solution  of  a  salt 
the  metal  of  which  forms  an  insoluble  chromate,  the  latter  is 
precipitated : 

2Ba(N03)2  +  K2Cr2O7  +  H2O  =  2BaCr04  +  2KNO3  +  2HNO3 

778.  Ammonium  dichromate,  (NH^C^O?,  is  used  in  the 
sensitive  material  employed  in  making  photographic  prints  by 
the  carbon  process.  Paper  is  coated  in  the  dark  with  a  solution 
of  gelatin  and  ammonium  dichromate,  in  which  lamp-black  or 
other  finely  divided  pigment  is  suspended,  and  allowed  to  dry. 
When  it  is  exposed  to  light  under  a  negative,  a  reaction  takes  place 
in  the  parts  of  the  sensitive  surface  illuminated;  the  dichromate  is 
reduced  to  chromium  trioxide,  which  renders  the  gelatin  in  con- 
tact with  it  insoluble  in  warm  water.  The  print  is  now  washed 
carefully  to  remove  the  part  of  the  sensitive  surface  which  has 
not  been  acted  upon  by  light,  and  the  portion  which  has  been 
developed  is  transferred  to  a  suitable  support. 

Powdered  ammonium  dichromate  burns  freely  with  a  flame 
when  it  is  ignited  by  means  of  a  piece  of  paper  which  has  been 
previously  dipped  into  a  solution  of  potassium  nitrate  and  dried. 
The  chromium  changes  in  valence  from  6  to  3,  the  oxygen  liberated 
burns  the  ammonia  to  nitrogen,  and  the  chromic  oxide  produced 
is  formed  as  a  light,  voluminous,  green  powder.  The  reaction 


CHROMIUM,  MOLYBDENUM,  TUNGSTEN,  AND  URANIUM     631 

which  is  a  very  striking  one  is  represented  by  the  following  equa- 
tion:   (NH4)2Cr2O7  =  N2  +  4H2O  +  Cr2O3. 

779.  The  Use  of  Chromic  Anhydride  and  Bichromates  as 
Oxidizing  Agents. — The  readiness  with  which  chromic  anhydride 
gives  up  a  part  of  its  oxygen  is  utilized  extensively  in  effecting 
processes  of  oxidation.  The  compound  chars  paper  and  ignites 
alcohol  at  ordinary  temperatures.  In  the  presence  of  acids  it 
is  an  active  oxidizing  agent  in  aqueous  solutions.  The  acid  facil- 
itates the  process,  since  it  unites  with  the  trioxide  formed  as  the 
result  of  the  reduction  of  the  anhydride: 

2CrO3  =  [Cr2O3]  +  3O 
[Cr2O3]  +  3H2SO4  =  Cr2(SO4)3  +  3H2O 


2CrO3  +  3H2SO4  =  O2(SO4)3  +  3H2O  +  30 

Three  oxygen  atoms  become  available  from  2  molecules  of  the 
anhydride  if  a  substance  is  present  that  can  be  oxidized.  For  most 
oxidations,  sodium  or  potassium  dichromate  and  sulphuric  acid 
are  commonly  used  because  the  salts  are  cheaper  than  the  anhy- 
dride. The  acid  first  liberates  the  anhydride  from  the  salt,  and 
then  the  reactions  indicated  above  take  place.  If  the  reaction 
K2Cr2O7  +  H2S04  =  K2SO4  +  H2O  +  2CrO3  is  added  to  these, 
the  combination  which  expresses  the  reaction  when  the  salt  is 
used  is  as  follows: 

K2Cr207  +  4H2SO4  =  K2SO4  +  Cr2(SO4)3  +  4H2O  +  [30] 

This  reaction  occurs  only  when  something  is  present  that  can  be 
oxidized,  and  the  fact  is  indicated  by  writing  the  formula  for  oxygen 
as  0;  the  equation  which  indicates  the  oxidation  of  ferrous  sul- 
phate to  ferric  sulphate  is  as  follows: 

2FeSO4  +  H2SO4  +  O  =  Fe2(SO4)3  +  H2O 

In  order  to  combine  this  equation  with  the  one  given  above,  it 
must  be  multiplied  by  3  so  that  the  amount  of  oxygen  available 
for  the  oxidation  is  equal  to  that  taken  up  by  the  substance 


632  INORGANIC  CHEMISTRY  FOR  COLLEGES 

oxidized.     When  this  has  been  done  and  the  partial  equations 
are  added  we  have  the  following  result : 

K2Cr2O7  +  4H2SO4  =  K2S04  +  Cr2(SO4)3  +  4H2O  +  [3O] 
6FeSO4  +  3H2S04  +  [3O]  =  3Fe2(SO4)3  +  3H2O 

K2Cr2O7  +  6FeSO4  +  7H2SO4 

=  K2SO4  +  Cr2(SO4)3  +3Fe2(SO4)3  +  7H2O 

The  reaction  which  takes  place  between  potassium  dichromate  and 
hydrochloric  acid  can  be  written  in  a  similar  way : 

K2Cr2O7  +  8HC1  =  2KC1  +  2CrCl3  +  4H2O  +  [3O] 
6HC1  +  [3O]  =  3H2O  +  3C12 

K2Cr2O7  +  14HG1  =  2KC1  +  2CrCl3  +  7H2O  +  3C12 

780.  Reactions  of  Oxidation  from  the  Point  of  View  of  Positive 
and  Negative  Valence. — When  mercuric  chloride,  HgCl2,  is 
treated  with  stannous  chloride,  SnCl2,  the  latter  is  oxidized  to 
stannic  chloride,  SnCl4,  and  the  former  reduced  to  mercury.  In 
mercuric  chloride  the  mercury  has  the  valence  +2,  because  it  is 
combined  with  2  atoms  of  a  univalent  negative  element;  in  the 
metallic  condition  its  valence  is  0.  The  change  in  valence  is 
—2,  because  2,  the  original  valence,  —2=0,  the  final  valence. 
The  valence  of  tin  in  stannous  chloride  is  +2  and  in  stannic  chlo- 
ride +4,  and,  accordingly,  as  the  result  of  the  reaction  the  valence 
of  tin  has  changed  by  +2,  for  +2,  the  original  valence,  +2  =  +4, 
the  final  valence.  The  mercury  "  loses  2  valencies  "  and  the  tin 
takes  these  up,  for  it  gains  2.  As  a  consequence,  1  molecule  of 
mercuric  chloride  will  react  with  1  molecule  of  stannous  chloride: 
HgCl2  +  SnCl2  =  Hg  +  SnCl4. 

Since  reactions  of  oxidation  of  this  type  involve  the  transfer 
of  an  element  or  a  radical  from  one  element  to  another,  the  loss  in 
valence  of  the  atoms  of  the  element  in  the  oxidizing  agent  must, 
evidently,  equal  numerically  the  gain  in  valence  of  the  atoms  of 
the  element  in  the  reducing  agent,  or,  expressed  in  other  words,  the 
two  changes  in  valence  must  be  equal  but  of  different  sign.  From 
the  point  of  view  of  +  and  —  valence  an  element  is  oxidized 


CHROMIUM,  MOLYBDENUM,  TUNGSTEN,  AND  URANIUM     633 

when  there  is  an  increase  in  its  positive  valence  (or  a  decrease  in 
its  negative  valence) ;  and  the  reverse  is  true  in  reduction. 

In  the  case  of  the  reduction  of  ferric  chloride,  Feds,  by  stan- 
nous  chloride,  the  iron  salt  is  reduced  to  ferrous  chloride,  FeCl2, 
and  the  change  in  valence  of  the  metal  is  —1,  for  +3  —  1  =  +2. 
The  change  of  valence  of  the  tin  in  stannous  chloride  is  +2,  con- 
sequently, 1  molecule  of  it  will  reduce  2  molecules  of  ferric  chloride: 
2FeCl3  +  SnCl2  =  2FeCl2  +  SnCl4.  • 

The  same  principle  applies  when  the  negative  valencies  are 
involved  in  oxidation  reactions.  When  ammonia  is  oxidized  to 
nitric  oxide  the  change  in  the  valence  of  nitrogen  is  from— 3  in 
ammonia  to  +2  in  nitric  oxide,  NO.  It  will  be  recalled  that  the 

<-H 
electronic  formulas  of  these  two  compounds  are  written  N  <—  H 

<-H 

and  N       0,  to  indicate  that  the  electrons  involved  pass  from  the 

more  positive  to  the  less  positive  element  (442);  and  it  will  also 
be  remembered  that  one  so-called  positive  valence  is  established 
on  an  element  for  each  electron  lost  and  one  negative  valence  for 
each  electron  gained.  The  total  change  in  valence  in  oxidizing 
ammonia  to  nitric  oxide  is,  therefore,  from  —3  to  +2  which  is 
+5,  for  —3  +  5  =  +2.  The  change  in  valence  of  the  element 
in  the  oxidizing  agent  will  be,  accordingly,  —5.  If  free  oxygen 
is  the  agent  employed,  it  changes  in  valence  from  0  to  —2  (H2O) 
and,  consequently,  2J  atoms  will  be  required  for  each  molecule 
of  ammonia  oxidized,  1NH3  to  2?  O,  or  expressed  as  a  molecular 
ratio,  4NH3  to  5O2. 

When  ammonia  is  treated  with  chlorine  it  is  oxidized  to  nitro- 
gen and  hydrochloric  acid  is  formed;  in  this  case  the  change  in 
valence  of  nitrogen  is  from  -3  to  0,  which  is  +3,  and  of  chlorine 
from  0  to  —1;  consequently,  for  each  molecule  of  ammonia  oxi- 
dized three  atoms  of  chlorine  will  be  required:  1NH3  to  3C1  or 
2NH3  to  3C12:  2NH3  +  3C12  =  N2  +  6HC1. 

781.  The  reactions  of  oxidation  involving  the  use  of  potassium 
dichromate  will  next  be  considered  from  this  point  of  view.  In 
the  salt  the  valence  of  chromium  is  6,  because  it  is  derived  from 
chromic  acid,  the  anhydride  of  which  is  CrO3.  Another  way  of 
determining  the  valence  of  chromium  is  to  consider  the  positive 


634  INORGANIC  CHEMISTRY  FOR  COLLEGES 

and  negative  valencies  of  the  atoms  of  which  the  salt  is  made  up. 
In  K2Cr2O?  these  are  2  potassium  atoms  each  having  the  valence 
+  1,  and  7  oxygen  atoms  each  having  the  valence  —2.  In  any 
compound  the  sum  of  the  negative  valencies  equals  the  sum 
of  the  positive  valencies,  or,  in  other  words,  the  sum  of  the  valen- 
cies is  numerically  equal  to  0  (442) .  Since  in  potassium  dichromate 
there  are  7  X  —  2  =  — 14  valencies  due  to  the  oxygen,  there  must 
be  +14  valencies;  of  these  the  two  potassium  atoms  furnish 
2  and  the  rest,  12,  must  be  furnished  by  the  two  chromium  atoms, 
which,  accordingly,  have  6  each. 

When  potassium  dichromate  acts  as  an  oxidizing  agent  in  the 
presence  of  acids,  the  chromium  is  reduced  to  the  valence  +3  and 
chromium  salts  are  formed:  the  change  in  valence  is  from  6  to  3 
or  —3,  and  since  each  molecule  of  the  salt  contains  2  atoms  of 
chromium  the  total  change  per  molecule  is  —6.  This  amount  of 
the  dichromate  will,  accordingly,  oxidize  the  amount  of  another 
substance  which  involves  a  change  of  +  6  in  valence.  If  hydro- 
chloric acid  is  to  be  oxidized  to  chlorine  the  change  in  valence 
of  chlorine  is  from  —1  to  0,  which  is  +1;  consequently,  1  molecule 
of  potassium  dichromate  will  oxidize  6  molecules  of  hydrochloric 
acid  to  6  atoms  of  chlorine.  More  of  the  acid  will  be  required 
than  this  amount,  however,  because  the  chromium  changes  to 
the  valence  3  in  which  it  is  a  base-forming  element  and,  con- 
sequently, is  converted  into  chromic  chloride,  CrCla.  Since 
2  atoms  of  the  element  are  involved,  6  molecules  of  hydrochloric 
acid  will  be  required  to  furnish  the  chlorine  for  this  purpose. 
And,  further,  the  potassium  in  the  salt  is  also  converted  into 
chloride  and  2  molecules  of  the  acid  will  be  needed  for  this  pur- 
pose, making  a  total  of  6  +  6  +  2  =  14  molecules  of  the  acid. 
The  equation  can  now  be  readily  written:  K^C^O?  +  14HC1  = 
2KC1  +  2CrCl3  +  3C12  +  7H2O.  It  should  be  noted  that  the 
oxygen  in  the  dichromate  is  converted  into  water,  the  hydro- 
gen for  the  purpose  coming  from  the  acid.  This  fact  makes  it 
possible  to  see  by  inspection  how  many  molecules  of  acid  are 
required.  In  this  case  there  are  the  7  oxygen  atoms,  and  the  14 
atoms  of  hydrogen  needed  are  furnished  by  14  molecules  of  hydro- 
chloric acid. 

782.  If  we  consider  next  the  oxidation  of  ferrous  sulphate, 
FeSO4,  to  ferric  sulphate,  Fe2 (864)3,  by  potassium  dichromate  in 


CHROMIUM,  MOLYBDENUM.  TUNGSTEN,  AND  URANIUM     635 

the  presence  of  sulphuric  acid,  the  total  change  in  valence  for  1 
molecule  of  the  dichromate  is  —6.  When  iron  changes  from  the 
ferrous  condition  in  which  the  valence  of  the  metal  is  2  to  the  ferric 
condition  in  which  it  is  3  the  change  in  valence  is  +1;  as  a  con- 
sequence, 1  molecule  of  potassium  dichromate  will  oxidize  6 
molecules  of  ferrous  sulphate.  The  number  of  molecules  of 
sulphuric  acid  required  is  determined  as  shown  above  by  the 
number  of  oxygen  atoms  in  the  oxidizing  agent;  the  14  atoms 
necessary  to  convert  7  atoms  of  oxygen  into  water  are  furnished 
by  7  molecules  of  sulphuric  acid.  The  completed  reaction  is, 
accordingly,  as  follows: 

K2Cr2O7  +  6FeS04  +  7H2S04 

=  K2S04  +  Cr2(S04)3  +3Fe2(S04)3  +  7H2O 

As  we  have  seen,  when  1  molecule  of  potassium  dichromate  acts 
as  an  oxidizing  agent  in  the  presence  of  an  acid,  the  valence  change 
is  —6.  When  oxygen  gas  acts  as  an  oxidizing  agent,  the  change  in 
valence  for  each  atom  is  —2,  for  the  element  has  this  valence  in 
its  compounds;  consequently,  1  molecule  of  the  dichromate  is 
equivalent  in  oxidizing  power  to  3  atoms  of  oxygen.  The  examina- 
tion of  reactions  of  oxidation  from  this  point  of  view  is  often  very 
helpful.  For  example,  in  oxidizing  alcohol,  C2H6O,  to  acetic 
acid,  C2H4O2,  2  atoms  of  oxygen  are  necessary — 

C2H6O  +  2O  =  C2H402  +  H2O 

As  a  consequence,  we  see  that  2  molecules  of  the  dichromate,  which 
are  equivalent  to  6  oxygen  atoms,  will  be  required  to  oxidize  3 
molecules  of  alcohol  which  require  6  of  oxygen  to  effect  their 
oxidation.  The  equation  for  the  reaction  is  as  follows: 

K2Cr2O7  +  3C2H6O  +  4H2S04 

=  K2S04  +  Cr2(S04)3  +  3C2H402  +  3H2O 

In  this  case  but  4  molecules  of  the  acid  are  required  because  only 
the  potassium  and  chromium  are  converted  into  sulphates. 

783.  Analytical  Reactions  of  Chromium. — Solutions  of  chromic 
salts  are  bluish-violet  or  green  in  color  (773)  and  yield,  when 
treated  with  a  soluble  hydroxide,  a  bluish-green  precipitate  of  the 
hydroxide,  which  is  insoluble  in  ammonia.  All  the  compounds  of 
chromium  when  fused  with  sodium  carbonate  and  potassium 


636  INORGANIC  CHEMISTRY  FOR  COLLEGES 

nitrate  are  converted  into  a  chromate  which  is  yellow;  when  fused 
with  borax  they  give  a  green  bead,  which  is  not  changed  in  color 
in  the  reducing  flame.  The  soluble  chromates  give  yellow  and 
the  dichromates  red  solutions;  they  are  reduced  by  hydrogen 
sulphide,  and  insoluble  chromic  hydroxide  and  sulphur  are  formed. 

784.  Molybdenum. — This  element  forms  a  number  of  oxides 
and  chlorides,  but  its  most  important  compound  is  molybdic  anhy- 
dride, MoOs,  from  which  a  large  number  of  salts  have  been  made. 
The  element  is  obtained  chiefly  from  molybdenite,  MoS2,  by  con- 
verting it,  by  roasting,  into  the  trioxide,  which  is  then  reduced  by 
heating  it  with  hydrogen.     Since  the  metal  melts  at  a  very  high 
temperature  it  cannot  be  worked  in  the  usual  way;    the  powder 
obtained  as  the  result  of  the  reduction  is  pressed,  and  the  ends 
of  the  blocks  connected  with  terminals  which  furnish  a  powerful 
electric  current.     The  heat  generated  by  the  passage  of  the  current 
causes  the  particles  to  coalesce.    Molybdenum  is  used  in  steel  and  is 
made  into  terminals  for  spark-plugs.     To  obtain  it  in  the  form  of 
wire,  the  metal  is  drawn  while  red-hot  through  dies  made  of  dia- 
mond set  in  steel. 

Ammonium  molybdate,  (NH^MoCU,  which  is  made  by  dis- 
solving molybdic  anhydride  in  ammonia,  is  used  in  the  laboratory 
in  testing  for  phosphates  and  in  their  quantitative  determination. 
When  a  phosphate  is  added  to  a  solution  of  the  reagent  in  dilute 
nitric  acid,  and  the  mixture  is  heated,  a  yellow  precipitate  of 
ammonium  phosphomolybdate  is  produced.  The  salt,  which 
has  the  formula  (NH4)3PO4,12MoO3,6H2O,  is  insoluble  in  dilute 
acids,  but  dissolves  in  alkalies. 

785.  Tungsten. — The   element   occurs   as   tungstates.     When 
these  are  fused  with  sodium  carbonate  they  are  converted  into 
sodium  tungstate,  Na2WC>4,  which  is  soluble  in  water.     Acids 
precipitate  from  the  solution  tungstic  acid,  EbWO^H^O,  which 
yields  tungsten  on  reduction  with  hydrogen.     The  metal  is  used 
extensively  in  making  high-speed  steel  (748)  and  as  a  filament  in 
electric-light  bulbs.     It  is  obtained  in  the  form  of  wire  in  the  way 
just  described  in  the  case  of  molybdenum.     Tungsten  melts  at  a 
higher  temperature  (3540°)  than  any  other  element,  and  the  dis- 
covery of  a  method  to  change  into  the  form  of  a  wire  the  powder 
formed  as  the  result  of  the  reduction  of  the  oxide,  revolutionized 
the  electric-light  industry.     Carbon  filaments  were  formerly  used 


CHROMIUM,  MOLYBDENUM,  TUNGSTEN,  AND  URANIUM     637 

in  electric  light  bulbs,  but  when  they  were  heated  to  a  very  high 
temperature  the  carbon  vaporized,  and  depositing  on  the  glass 
darkened  it;  in  a  short  time  the  filament  became  disintegrated. 
On  account  of  its  high  melting-point  and  low  volatility,  tungsten 
can  be  heated  to  a  much  higher  temperature  than  carbon  without 
acting  in  this  way,  and  since  the  quantity  of  heat  converted  into 
light  increases  very  rapidly  with  rise  in  temperature  (455),  the 
efficiency  of  a  tungsten  lamp  is  much  greater  than  that  of  one  in 
which  the  filament  is  made  of  carbon.  A  lamp  of  the  latter  type 
requires  3.25  watts  per  candle-power,  whereas  a  tungsten  lamp 
requires  only  1.25  watts  to  produce  the  same  amount  of  light. 

786.  Uranium. — The  element  occurs  in  small  quantities  in 
several  minerals  which  are  complex  in  composition;  in  pitchblende 
it  is  present  as  an  oxide,  UsCU,  associated  with  a  large  number  of 
other  elements;  carnotite,  which  has  the  composition  K2O,2UOs,- 
V2Os,3H2O,  occurs  in  Colorado  and  is  used  as  a  source  of  radium 
which  is  found  in  minute  amounts  in  uranium  ores.  The  element 
forms  a  large  number  of  compounds,  the  most  important  of  which 
are  derived  from  the  oxide  of  the  composition  UOs,  which  is  an  acid 
anhydride  and  yields  salts  analogous  in  composition  to  the  chro- 
mates  and  dichromates.  Sodium  uranate  has  the  formula  Na2UC>4, 
and  sodium  diuranate,  the  formula  Na2U2O7,7H2O.  The  latter 
salt  is  used  in  coloring  uranium  glass,  which  exhibits  the  phe- 
nomenon of  fluorescence,  and  is  yellow  or  green  depending  on  the 
way  the  glass  is  viewed.  Uranium  trioxide  has  weak  base-forming 
properties;  it  forms  basic  salts  of  which  uranyl  nitrate,  UO2(NO3)2, 
7H2O  and  uranyl  sulphate,  UC^SCU,  3JH20,  are  examples.  The 
radical  which  occurs  in  these  salts,  UC>2++,  is  called  uranyl;  the 
name  recalls  that  of  sulphuryl  chloride,  SO2C12  (290).  The 
uranyl  salts  are  yellow  and  their  solutions  fluorescent,  being  yellow 
by  reflected  and  green  by  transmitted  light.  Uranium  and  its 
compounds  are  of  especial  interest  on  account  of  the  fact  that  they 
exhibit  the  phenomenon  of  radioactivity  (799). 

EXERCISES 

1.  Compare  the  chemical  behavior  of  (a)  A1(OH)3,  Fe(OH)3,  andCr(OH)3, 
(6)  Cr(OH)2  and  Fe(OH)2. 

2.  Write  equations  for  all  the  reactions  involved  in  making  K2Cr2O7 
from  chrome-iron  ore.      State  why  the  process  serves  to  separate  the  iron 
from  the  chromium. 


638  INORGANIC  CHEMISTRY  FOR  COLLEGES 

3.  Write  ionic  equations  for  the  reactions  which  occur  when  a  solution 
of  Ba(NO3)2  is  treated  with  one  of  K2Cr2O7. 

4.  Calculate  the  valence  of  the  acid-forming  element  in  the  following 
compounds,    using  -f-   and  —    valencies:     (a)  CrPO4,     (6)  K2Mn5On,     (c) 
KH2PO2,    (d)  CaSO3,    (e)  (NH4)2Cr3O,0. 

5.  Under  certain   conditions  alcohol,   C2H60,   is  oxidized  to  aldehyde, 
C2H4O.     Write  the  equation  for  the  reaction  which  takes  place  when  the 
oxidation  is  effected  by  means  of  K2Cr2O7  in  the  presence  of  H2SO4. 

6.  Write  an  equation  for  the  oxidation  of  H2S  by  K2Cr2O7  as  the  result 
of  which  S  and  Cr(OH)3  are  formed. 

7.  How  could  you  distinguish  from  each  other  the  salts  of  the  following 
metals:    (a)  Cr  and  Ni,    (6)  Cr  and  Fe,    (c)  Cr  and  Co,    (d)    How  could 
you  separate  Fe(OH)3  from  Cr(OH)3? 

8.  How  could  you  obtain  pure  metallic  iron  and  pure  metallic  chromium 
starting  with  a  mixture  of  the  sulphates  of  the  metals? 


CHAPTER  XLIV 

MANGANESE 

787.  We  have  just  seen  from  a  study  of  the  chemistry  of 
chromium  that  the  valence  of  an  element  is  the  factor  which 
largely  determines  the  chemical  behavior  of  its  compounds.     This 
important  conclusion  is  even  more  strikingly  illustrated  in  the 
case  of  manganese,  which  forms  compounds  in  which  the  element 
shows  the  valencies  2,  3,  4,  6  and  7.     In  the  elementary  state 
manganese  has  the  properties  characteristic  of  metals  and  resem- 
bles closely  iron.     When  it  has  the  valence  2  it  forms  salts  like 
those  of  magnesium  and  is  a  relatively  strong  base-forming  ele- 
ment.    With  the  valence  3,  it  functions  as  a  weak  base-forming 
element  and  yields  salts  which  resemble  those  of  aluminium  al- 
though they  are  much  less  stable  and  are  more  readily  hydrolyzed. 
With  increase  in  the  valence  of  the  element  to  4,  compounds  are 
produced  which  are  like  those  of  lead  when  it  shows  this  valence; 
the  oxide  of  the  composition  MnO2  has  very  weak  acid-forming 
properties,  and  the   corresponding   chloride,   MnCU,   is  unstable 
and  like  lead  tetrachloride  breaks  down  into  the  dichloride  and 
chlorine. 

The  element  becomes  strictly  acid-forming  in  character  when 
it  has  the  valence  6;  it  forms  salts,  of  which  potassium  manga- 
nate,  K^MnCU,  is  an  example,  that  are  isomorphous  with  the 
sulphates.  Its  highest  valence  is  7,  a  fact  which  is  in  accord 
with  its  position  in  the  seventh  group  in  the  periodic  classification 
of  the  elements.  With  this  valence  manganese  functions  as  a 
strong  acid-forming  element  and  yields  salts  called  permanganates, 
which  resemble  the  perchlorates  in  both  physical  and  chemical 
properties;  potassium  permanganate,  KMnO4,  and  potassium 
perchlorate,  KC1O4,  are  examples  of  such  salts. 

788.  Preparation   and   Properties   of   Manganese. — The  ele- 
ment occurs  chiefly  as  oxides  and  hydrated  oxides.     Braunite, 

,  manganite,  MnO  (OH)  or  M^Oa ,  H2O,  hausmannite, 
639 


640  INORGANIC  CHEMISTRY  FOR  COLLEGES 

and  manganese  spar,  MnCOs,  are  important  minerals.  The 
chief  ore  of  the  element  is  pyrolusite,  Mn(>2.  Manganese  may 
be  obtained  by  reducing  an  oxide  of  the  metal  with  aluminium, 
but  the  large  amounts  required  in  making  iron  alloys  are  made 
by  reducing  a  mixture  of  an  iron  ore  and  a  manganese  ore  in  a 
blast  furnace.  Spiegeleisen  and  ferromanganese  are  made  in  this 
way;  the  former  contains  10  to  20  per  cent  manganese,  4  to  5 
per  cent  carbon,  and  the  rest  is  iron;  the  latter  contains  20  to 
85  per  cent  manganese  and  6  to  7  per  cent  carbon. 

Manganese  resembles  iron  in  appearance  but  its  luster  has  a 
reddish  tinge;  it  is  hard  and  brittle,  has  the  specific  gravity  7.2 
and  melts  at  1260°.  It  oxidizes  in  the  air,  decomposes  steam  at 
red  heat,  and  dissolves  in  acids  to  form  manganous  salts.  When 
the  element  or  any  of  its  oxides  is  heated  in  the  air  the  final  product 
has  the  composition  MnaO^ 

789.  Manganous  Compounds . — Manganous  hydroxide, 
Mn(OH)2,  is  precipitated  when  a  base  is  added  to  a  solution  of  a 
manganous  salt;  when  pure  it  is  white,  but  it  is  soon  oxidized  by 
the  air  more  or  less  completely  to  manganic  hydroxide,  Mn(OH)3, 
which  is  brown;  the  reaction  is  like  that  which  takes  place  in 
the  case  of  ferrous  hydroxide.  The  hydroxide,  like  ferrous  and 
magnesium  hydroxides,  is  not  precipitated  by  ammonium  hydrox- 
ide in  the  presence  of  ammonium  salts  (589). 

Manganous  chloride,  MnCl2,4H2O,  is  obtained  by  treating 
manganese  dioxide  with  hydrochloric  acid;  the  salt  crystallizes 
from  water,  is  pale  pink  in  color,  and  is  stable  in  the  air.  Man- 
ganous sulphate,  MnS(>4,  forms  a  number  of  hydrates;  the  one 
obtained  by  evaporation  of  a  solution  at  room  temperature  is  a 
pentahydrate,  which  is  isomorphous  with  copper  sulphate. 
Manganous  carbonate,  MnCOs,  is  white  when  pure,  but  is  usually 
colored  brown  owing  to  the  presence  of  manganic  hydroxide  which 
is  formed  as  the  result  of  the  oxidation  of  the  small  amount  of 
manganous  hydroxide  produced  as  the  result  of  the  partial 
hydrolysis  of  the  carbonate.  It  is  decomposed  at  a  compara- 
tively low  temperature  by  heat,  and  for  this  reason  is  used  under 
certain  conditions  in  the  laboratory  as  a  source  of  carbon  dioxide. 
Manganous  sulphide,  MnS,  is  formed  by  the  action  of  ammonium 
sulphide  on  a  manganous  salt;  it  is  a  flesh-colored  amorphous 
precipitate,  which,  like  ferrous  sulphide,  is  soluble  in  dilute  acids. 


MANGANESE  641 

All  the  soluble  manganous  salts  have  a  pink  color;  those 
derived  from  the  strong  acids  are  but  slightly  hydrolyzed  and 
are  stable  in  the  air,  and  in  the  latter  respect  differ  from  the 
salts  of  stannous  tin  and  ferrous  iron. 

790.  Manganic  Compounds. — Manganic  hydroxide,  Mn(OH)s, 
is  formed  as  the  result  of  the  hydrolysis  of  manganic  sulphate, 
Mn2(SO4)3,   and  by  the  oxidation  of  the  solution  formed  by 
dissolving  manganous  hydroxide  in  solutions  of  ammonium  salts. 
It  is  brown-black  in  color  and  dissolves  in  hydrochloric  acid  to 
form  manganic  chloride,  MnCls,  which  readily  decomposes  into 
manganous    chloride    and    chlorine.     Like   other   unstable   com- 
pounds of  this  kind,  it  forms  double  salts  which  are  stable;    for 
example,  a  compound  of  the  formula  MnCls, 2KC1  can  be  pre- 
pared by  dissolving  manganese  dioxide  in  a  solution  containing 
hydrochloric  acid  and  potassium  chloride. 

The  marked  difference  between  iron  and  manganese  is  that  in 
the  case  of  the  former  the  salts  containing  the  element  in  the 
trivalent  condition  are  stable  in  the  air,  whereas  in  the  case  of 
manganese  salts,  the  salts  of  the  metal  in  the  bivalent  condition 
are  stable.  In  the  trivalent  condition  manganese  is  a  less  active 
base-forming  element  than  iron,  and  its  salts  are  more  com- 
pletely hydrolyzed  than  those  of  ferric  iron. 

791.  Manganites. — Manganese    dioxide,    MnO2,    like    silicon 
dioxide,  SiO2,  acts  as  an  acid  anhydride  when  fused  with  alkalies; 
like  other  anhydrides  of  this  class,  in  which  acid-forming  proper- 
ties are  developed  to  but  a  slight  degree,  it  yields  salts  that  are 
complex  in  composition  and  contain  a  number  of  molecules  of  the 
anhydride  to  one  molecule  of  the  basic  oxide.     There  are  many 
silicates,  chromates,  tungstates,  and  molybdates  of  this  type  that 
are  derived  from  anhydrides  which  are  very  weakly  acidic. 

When  manganese  dioxide  is  fused  with  potassium  hydroxide,  a 
number  of  manganites  are  formed,  one  of  which  has  the  formula 
K2O,5MnO2.  When  calcium  hydroxide  is  added  to  a  solution  of 
a  manganese  salt  and  air  is  blown  through  the  mixture,  the  man- 
ganous hydroxide  precipitated  is,  in  part,  oxidized  in  the  pres- 
ence of  the  base,  and  manganites  of  the  formulas  CaO,MnO2 
and  CaO,2MnO2  are  produced.  A  part  of  the  manganese  remains 
in  the  bivalent  condition  and  unites  with  some  of  the  dioxide  to 
form  manganous  manganite,  MnO,MnO2.  These  reactions  are 


642  INORGANIC  CHEMISTRY  FOR  COLLEGES 

utilized  in  recovering  the  manganese  used  in  the  process  for  the 
manufacture  of  chlorine  from  manganese  dioxide  and  hydro- 
chloric acid.  The  so-called  Weldon  "  mud  "  obtained  as  out- 
lined above  is  treated  with  hydrochloric  acid,  and  chlorine  is 
obtained. 

792.  Manganates. — When  any  compound  of  manganese  is 
fused  with  potassium  carbonate  and  potassium  nitrate,  the  melt 
becomes  green  in  color  and  potassium  manganate,  K2MnO4,  is 
obtained:  MnO2  +  K2CO3  +  O  =  K2MnO4  +  C02.  The  salt 
crystallizes  from  water  in  anhydrous  greenish-black  crystals, 
which  give  a  deep  green  solution  when  dissolved  in  water. 

Manganese  can  have  the  valence  6  only  when  it  is  in  com- 
bination with  a  base-forming  element.  If  an  acid  is  added  to  a 
solution  of  a  manganate,  the  manganic  acid,  H2MnO4,  liberated 
decomposes  at  once  into  permanganic  acid,  HMnO4,  and  man- 
ganese dioxide.  The  change  can  be  readily  observed  because 
the  manganates  are  green  and  permanganic  acid  and  its  salts  are 
red.  When  a  manganate  is  dissolved  in  water  it  hydrolyzes  to 
some  extent: 

K2MnO4  +  2H20  ^±  2KOH  +  H2MnO4 
The  manganic  acid  formed  then  decomposes: 

3H2Mn04  =  MnO2  +  2HMnO4  +  2H20 

The  permanganic  acid  next  interacts  with  a  part  of  the  alkali 
formed  as  the  result  of  the  hydrolysis  of  the  manganate : 

KOH  +  HMnO4  =  KMnO4  +  H2O 

The  combination  of  these  partial  equations  leads  to  the  following 
equation: 

3K2MnO4  +  2H2O  =  2KMnO4  +  4KOH  +  MnO2 

Since  potassium  manganate  is  stable  in  the  presence  of  alkalies, 
its  hydrolysis  takes  place  completely  only  if  an  acid  is  added  to 
neutralize  the  hydroxide  formed  as  the  result  of  the  hydrolysis. 
It  is  for  this  reason  that  carbon  dioxide  is  passed  into  the  solution 
in  effecting  the  change  of  manganate  to  permanganate,  and  when 


MANGANESE  643 

the  equation  for  this  reaction  is  combined  with  the  one  just  given 
we  arrive  at  the  final  equation  for  the  reaction : 

3K2MnO4  +  2CO2  =  2KMnO4  +  2K2CO3  +  MnO2 

793.  Permanganates. — Potassium  permanganate  was  formerly 
prepared  by  making  use  of  the  reactions  just  discussed,  but  since 
the  process  involved  the  conversion  of  one-third  of  the  man- 
ganese into  the  dioxide,  the  salt  is  now  manufactured  by  oxidizing 
a  solution  of  potassium  manganate  by  means  of  chlorine : 

2K2MnO4  +  C12  =  2KMnO4  +  2KC1 

It  is  also  prepared  by  direct  oxidation  in  an  electrolytic  cell: 
2K2MnO4  +  O  +  H2O  =  2KMnO4  +  2KOH 

The  permanganate  separates  as  crystals  near  the  anode.  The 
hydroxide  formed  collects  at  the  cathode,  and  is  used  in  converting 
more  manganese  dioxide  into  manganate  by  the  fusion  process. 

Potassium  permanganate  dissolves  in  cold  water  (1  part  in  16) 
and  forms  purple  crystals.  It  is  an  active  oxidizing  agent,  which 
can  ignite  certain  organic  compounds;  when  a  drop  of  glycerine 
is  placed  on  the  powdered  salt,  the  former  burns  with  a  blue  flame. 
Potassium  permanganate  is  used  as  a  disinfectant,  in  bleaching 
and  dyeing,  and  in  medicine.  Sodium  permanganate  was  used 
during  the  recent  war  in  gas  masks  as  a  protection  against  arsine 
and  other  oxidizable  gases,  and  was  incorporated  into  granules 
containing  sodium  hydroxide,  lime,  and  cement.  The  bases  were 
used  to  decompose  halogen  compounds  and  the  cement  to  render 
the  mass  porous. 

Permanganic  acid  can  be  obtained  in  the  form  of  hydrated 
crystals  by  adding  sulphuric  acid  to  a  solution  of  barium  perman- 
ganate, filtering  off  the  barium  sulphate,  and  evaporating  the  fil- 
trate at  a  low  temperature.  The  acid  decomposes  at  32°  into 
manganese  dioxide,  water,  and  oxygen.  It  has  also  been  prepared 
by  electrolyzing  a  solution  of  potassium  permanganate.  The 
"acid  collects  at  the  anode,  which  is  placed  in  a  porous  cup,  and 
potassium  hydroxide  at  the  cathode. 

When  potassium  permanganate  is  cautiously  treated  with 
concentrated  sulphuric  acid,  an  oil,  which  gives  a  violet  vapor,  is 


644  INORGANIC  CHEMISTRY  FOR  COLLEGES 

formed;  it  has  the  composition  Mi^O?,  and  is  the  anhydride  of 
permanganic  acid.  The  oil  is  explosive  and  sets  fire  to  paper  and 
other  organic  substances. 

794.  Potassium  Permanganate  as  an  Oxidizing  Agent. — The 
fact  that  potassium  permanganate  is  such  an  active  oxidizing 
agent,  makes  it  a  very  valuable  reagent  in  volumetric  quantitative 
analysis.  The  solution  of  the  salt  is  stable  provided  it  is  kept 
free  from  any  oxides  of  manganese,  which  catalyze  its  reduction 
to  manganese  dioxide.  Since  it  slowly  oxidizes  paper  and  other 
organic  substances,  it  must  be  protected  from  these  as  well  as 
from  the  dust  of  the  air,  which  contains  organic  material.  No 
indicator  is  needed  when  we  are  using  the  solution  as  an  oxidizing 
agent,  for  its  purple  color  disappears  when  it  has  been  reduced. 

When  potassium  permanganate  oxidizes  in  the  presence  of 
acids  it  is  reduced  to  a  manganous  salt  and  the  change  in  valence 
is  from  +7  to  +2.  If  ferrous  sulphate  is  being  oxidized,  for 
example,  to  ferric  sulphate,  the  change  in  valence  of  the  iron  is 
from  +2  to  +3;  as  a  consequence,  1  molecule  of  the  perman- 
ganate will  oxidize  5  molecules  of  ferrous  sulphate. 

When  the  oxidation  occurs  in  neutral  or  alkaline  solutions  the 
permanganate  is  reduced  only  to  manganese  dioxide  and  the 
valence  change  is  from  +7  to  +4.  It  is  important  to  remember 
that  in  acid  solution  the  salt  furnishes  oxygen  equivalent  to  5 
"  valencies  "  or  equivalents  (2J  atoms  of  oxygen),  and  in  neutral 
or  alkaline  solutions,  3  equivalents.  With  these  facts  in  mind 
it  can  be  seen  readily  that  1  molecule  of  potassium  permanganate 
will  oxidize  in  acid  solution  5  molecules  of  a  chromous"  salt  to  a 
chromic111  salt,  2J  molecules  of  a  stannous11  salt  to  a  stanniciv 
salt,  2|  molecules  of  sulphuriv  dioxide  to  sulphuricvi  acid,  5 
molecules  of  hydriodic"1  acid  to  iodine0,  2J  molecules  of  nitrous111 
acid  to  nitricv  acid,  etc. 

The  equations  for  the  reactions  can  be  written  by  applying  the 
method  described  at  length  when  chromic  acid  was  considered 
(779).  The  case  of  the  oxidation  of  ferrous  sulphate  in  the 
presence  of  sulphuric  acid  may  be  taken  as  an  example.  One 
molecule  of  the  permanganate  will  oxidize  5  molecules  of  the  salt, 
since  the  change  in  valence  of  the  manganese  is  5  and  that  of  the 
iron  salt  is  1.  Since  ferric  sulphate  contains  2  atoms  of  iron,  we 
must  take  an  even  number  of  molecules  of  ferrous  sulphate — 10 


MANGANESE  645 

molecules  of  the  latter  will  make  5  of  ferric  sulphate.  We  require 
for  this  amount,  evidently,  2  molecules  of  potassium  permanganate. 
These  contain  8  oxygen  atoms  which  are  converted  into  water  in 
the  reaction,  and,  consequently,  8  molecules  of  sulphuric  acid 
are  necessary  to  furnish  the  required  16  hydrogen  atoms.  The 
completed  reaction  is  then  as  follows: 

2KMnO4  +  10FeSO4  +  8H2SO4 

=  K2SO4  +  2MnSO4  +  5Fe2(SO4)3  +  8H2O 

795.  Analytical  Reactions  of  Manganese. — Compounds  of  the 
element  are  converted  into  a  green  manganate  when  fused  with 
potassium  carbonate  and  potassium  nitrate.  When  treated 
with  ammonium  sulphide  manganous  salts  give  a  flesh-colored 
precipitate  of  manganous  sulphide,  which  dissolves  in  dilute  acids. 
Manganates  are  green  and  permanganates  are  violet;  they 
are  converted  by  ferrous  salts,  zinc  and  sulphuric  acid,  and  by 
other  reducing  agents  into  manganous  salts,  which  are  very  light 
pink  or  practically  colorless  in  dilute  solutions. 

EXERCISES 

1.  How  could  Mn2O3  be  distinguished  from  Fe?O3  through  the  use  of 
HC1? 

2.  Complete    and    balance     the     following:       (a)  FeSO4-f  K2MnO4-f- 
H2SO4  =      ,    (b)  SnCl2  -f  K2MnO4  +  HC1  =      ,    (c)  KMnO4  +  HC1  =       , 
(d)  KMnO4  +  S02  +  H.SO4  =     ,    (e)  KMnO4  +  KNO2  +  H2SO4  =       . 

3.  The  strength  of  a  solution  of  KMnO4  can  be  determined  by  titratimg 
it  with    Na2S2O3  after  treatment  with  KI  and  dilute  sulphuric  acid.     Write 
equations  for  all  the  reactions  involved. 

4.  Write  equations  for  the  reactions  that  take  place  when  an  alloy  of 
Fe,  Mn,  and  Cr  is  dissolved  in  hot  aqua  regia. 

5.  Explain  why  Fe  can  be  separated  by  the  basic  acetate  method  (756), 
from  Mn  and  not  from  Cr. 

6.  (a)  Write  an  equation  for  the  reaction  by  which  alcohol,  C2H6O,  is 
oxidized  to  CO2  +  H2O  by  KMnO4  in  the  presence  of  H2SO4.     (6)  Write 
an  equation  for  the  oxidation  if  no  acid  is  present. 


CHAPTER  XLV 
RADIOACTIVITY.     THE  STRUCTURE  OF  ATOMS 

796.  The  discovery  of  X-rays  and  of  radium  (652)  brought 
to  light  a  new  type  of  energy  the  study  of  which  has  resulted  in 
extending   to   a  marked   degree   our   knowledge   of  matter  and 
its  transformation.     It  is  possible  to  count  individual  atoms  and 
determine  their  weight,  and  the  facts  which  have  been  discovered 
lead  to  the  reasonable  conclusion  that  the  atoms  of  the  elements 
are  made  up  of  simpler  units,  which  appear  to  be  hydrogen  and 
helium.     Before  the  facts  upon  which  these  views  are  based  can 
be  given,  some  of  the  phenomena  concerned  must  be  described 
briefly. 

797.  X-Rays. — A  tube  to  produce  X-rays  is  made  by  sealing 
into  a  glass  bulb  two  metallic  conductors,  and  exhausting  the 
bulb  to  the  highest  vacuum  obtainable.     When  an  electric  cur- 
rent from  an  induction  coil  is  sent  through  the  tube,  a  stream 
of  electrons,  which  are  charges  of  negative  electricity,  is  given 
off  from  the  cathode.     When  these  fall  upon  any  substance  placed 
opposite  the  cathode  (the  anti-cathode),  vibrations,  called  X-rays, 
are  produced  which  are  of  the  same  nature  as  those  of  visible 
light.     They  differ  from  the  latter  only  in  wave-length.     The 
wave-lengths  of  visible  light  are  of  the  order  of  10~5  cm.,  while 
those  of  the  X-rays  are  about  one-thousandth  as  large  (10~8  cm.). 
X-rays  will  pass  through  materials  through  which  the  longer 
light  rays  will  not  pass;   it  is  this  property  of  the  rays  which  is 
utilized  in  X-ray  photography. 

798.  Atomic  Numbers. — When  white  light  falls  upon  a  "  grat- 
ing "  made  by  ruling  lines  on  a  polished  metallic  surface,  the  light 
is  broken  up  into  its  constituents  and  a  spectrum  is  formed. 
In  order  to  obtain  this  effect,  the  distances  between  the  unscratched 
surfaces  of  the  grating  which  reflect  light  must  be  of  the  order 
of  the  wave-length  of  light.     Gratings  have  been  made  by  ruling 

646 


RADIOACTIVITY.     THE  STRUCTURE  OF  ATOMS         647 

200,000  lines  to  the  inch.  X-Rays  of  different  wave-lengths  are 
separated  in  a  similar  way,  when  they  fall  on  a  surface  which  is 
so  constituted  that  alternate  parts  will  reflect  the  waves.  It 
is  impossible  to  make  a  grating  of  this  kind  mechanically,  because 
the  distances  between  the  reflecting  surfaces  must  be  of  the 
order  of  the  wave-lengths  of  the  rays.  It  was  discovered  by 
Laue,  in  1912,  that  when  X-rays  fall  upon  a  crystal  of  salt,  the 
surface  of  the  crystal  acts  as  a  grating  as  a  result  of  the  orderly 
arrangement  of  its  atoms,  which  serve  as  the  reflecting  surface. 
In  this  way  it  is  possible  to  determine  the  wave  lengths  of  the 
X-rays  given  off  from  any  substance  when  it  is  the  anti-cathode 
in  an  X-ray  tube.  The  spectrum  produced  is  not  visible,  as  it 
is  in  the  case  of  light,  but  it  can  be  photographed  in  the  ordinary 
way,  because  X-rays  act  upon  a  photographic  plate. 

Moseley  in  1914  examined  the  X-rays  produced  when  different 
elements  served  as  the  anti-cathode  in  an  X-ray  tube.  He  found 
that  each  substance  produced  waves  of  a  definite  wave-length, 
which  are  characteristic  of  the  element.  And  what  was  more 
striking,  he  found  that  there  was  a  direct  relation  between  the 
atomic  weight  of  the  element  and  the  wave-length  of  the  X-rays 
produced.  As  the  atomic  weight  increased,  the  wave-length 
increased.  When  the  elements  were  arranged  according  to  the 
wave-lengths  of  the  X-rays  they  gave  off,  they  followed  in  the 
order  in  which  they  occur  in  the  periodic  classification.  The 
exceptions  which  are  met  with  when  the  elements  are  arranged 
according  to  their  atomic  weights,  disappeared,  however.  It 
will  be  recalled  that  the  atomic  weight  of  tellurium  places  it  in 
the  periodic  classification  in  the  halogen  family,  and  that  of  iodine 
in  the  sulphur  family,  and  the  order  in  which  iron,  cobalt,  and 
nickel  fall  is  anomalous.  When  these  elements  are  arranged 
according  to  the  wave-lengths  of  their  characteristic  X-rays,  they 
fall  in  the  order  demanded  by  their  chemical  properties.  It  is 
thus  seen  that  the  atomic  weight  of  an  element,  which  is  a  measure 
of  its  mass,  is  not  its  most  fundamental  property.  The  X-rays 
emitted  by  an  atom  when  it  is  struck  by  an  electron,  are  produced 
as  the  result  of  the  vibrations  set  up  in  the  atom.  These  are 
determined,  in  all  probability,  by  the  electrical  charges  on  the 
atom. 

In  classifying  the  elements  according  to  the  wave-lengths  of 


648  INORGANIC  CHEMISTRY  FOR  COLLEGES 

the  X-rays  they  emit,  a  number  is  assigned  to  each  element  accord- 
ing to  its  position  in  the  series.  Hydrogen  is  numbered  1,  helium 
2,  lithium  3,  etc.  The  number  assigned  to  each  element  is  called 
its  atomic  number.  The  number  of  each  element  is  given  in  the 
table  illustrating  the  periodic  classification,  which  is  printed  on 
the  page  facing  the  back  cover  of  this  book.  It  will  be  seen  that 
a  more  exact  statement  of  the  properties  of  elements  than  that 
enunciated  by  Mendelejeff,  can  now  be  given;  it  is — the  prop- 
erties of  the  elements  are  a  periodic  function  of  their  atomic  num- 
bers. Mendelejeff 's  generalization  was  based  on  the  mass  of  the 
atoms,  whereas  the  newer  classification  is  based  on  energy  rela- 
tionships, which  appear  to  be  more  fundamental. 

An  interpretation  of  the  atomic  numbers  leads  to  the  con- 
clusion that  there  are  92  elements,  since  92  is  the  atomic  number 
of  uranium,  and  we  have  no  evidence  that  heavier  elements  exist. 
Of  these,  87  have  been  discovered.  The  places  left  vacant  in  the 
table  and  indicated  by  dots,  .  . .  . ,  are  those  into  which  the 
five  undiscovered  elements  should  be  put;  their  atomic  numbers 
are  43,  61,  75,  85,  and  87,  respectively. 

799.  Radioactivity. — The  salts  of  radium  constantly  emit 
three  kinds  of  "  rays  "  which  differ  markedly  from  one  another. 
The  so-called  a-rays  have  been  shown  to  be  made  up  of  positively 
charged  helium  atoms,  which  are  produced  as  the  result  of  the  loss 
of  two  electrons  by  each  helium  atom.  It  will  be  recalled  that 
an  electron  is  a  negative  charge  of  electricity,  which  has  an  appar- 
ent mass  equal  to  about  18100  that  of  a  hydrogen  atom.  The 
charged  helium  atoms  are  called  a-particles;  they  each  carry 
the  charge  equal  to  that  on  a  bivalent  metallic  ion,  and  when  they 
lose  this  charge  the  helium  has  the  properties  of  the  gas  extracted 
from  natural  gas  and  other  sources  (331). 

The  /3-rays  consist  of  electrons  which  move  with  about  the 
same  velocity  as  that  of  light.  The  7-rays  are  X-rays  produced 
as  the  result  of  the  bombardment  of  the  atoms  of  radium  by  the 
electrons  which  are  given  off  by  the  element. 

During  the  disintegration  of  radium  a  gas  is  produced  which 
was  originally  called  radium  emanation;  it  was  shown  to  be  an 
element  that  resembled  argon  in  properties,  and  was  later  named 
niton.  A  determination  of  the  density  of  the  gas  gave  results 
that  led  to  the  atomic  weight  222.4.  The  atomic  weight  of  niton 


RADIOACTIVITY.     THE  STRUCTURE  OF  ATOMS         649 

places  it  in  group  0  of  the  periodic  classification  of  the  elements; 
and  this  position  is  in  accord  with  its  properties. 

The  fact  that  uranium  is  radioactive  led  to  the  view  that 
the  radium  which  is  associated  with  it  was  probably  produced 
from  uranium,  and  an  investigation  of  the  subject  showed  that 
the  view  was  correct.  A  sample  of  a  uranium  salt  which  had 
been  freed  from  radium  was  found  after  about  three  weeks  to 
contain  the  latter,  and  in  about  one  year  the  relation  between  the 
amounts  of  the  two  elements  present  became  constant  (1  gram  of 
uranium  to  3.2X10"7  grams  of  radium).  This  relationship 
makes  clear  why  such  small  quantities  of  radium  are  obtained 
from  pitchblende  (652);  from  1  ton  of  the  mineral  about  0.2 
gram  of  radium  can  be  extracted. 

800.  Radium  is  not  the  first  product  formed  from 
uranium;  the  latter  decomposes  into  helium  and  an  element 
called  uranium-Xi,  which  through  the  loss  of  electrons  only 
passes  into  uranium-X2;  this  gives  U2,  which  yields  helium  and 
ionium,  and  the  latter  then  changes  on  disintegration  into  helium 
and  radium.  The  decomposition  of  radium,  as  we  have  seen, 
produces  helium  and  niton.  The  radio-active  transformations 
continue  through  Ra-A,  Ra-B,  Ra-C,  Ra-Ci,  Ra-D,  Ra-E 
Ra-F,  and,  finally,  lead  is  produced.  When  helium,  which  has 
the  atomic  weight  4,  is  given  off  in  the  transformations,  the 
atomic  weight  of  the  element  formed  is  4  less  than  that  of  the 
element  from  which  it  was  produced.  In  the  complete  change 
from  uranium,  which  has  the  atomic  weight  238.2,  to  lead,  there 
are  eight  radioactive  transformations  which  produce  helium;  as  a 
consequence,  the  atomic  weight  of  lead  should  be  8  X  4  =  32 
less  than  the  atomic  weight  of  uranium,  that  is,  206.2.  Since 
the  atomic  weight  of  lead  is  207.2,  it  seemed  of  importance  to 
study  the  lead  that  was  associated  with  uranium  in  its  ores,  for 
the  metal  formed  under  these  conditions  should  have  a  lower 
atomic  weight  than  207.2,  if  the  view  is  correct  that  it  is  the  final 
disintegration  product  of  uranium. 

Lead  was  isolated  from  uranium  ores  and  after  it  had  been 
freed  from  all  other  elements  its  atomic  weight  was  determined. 
The  metal  obtained  from  different  radioactive  sources  gave  values 
for  the  atomic  weight  which  varied  from  206.4  to  206.8.  The 
compounds  prepared  from  the  lead  which  has  the  atomic  weight 


650  INORGANIC  CHEMISTRY  FOR  COLLEGES 

206.4  possess  the  identical  physical  and  chemical  properties  ex- 
hibited by  the  analogous  compounds  prepared  from  ordinary  lead, 
which  has  the  atomic  weight  207.2.  These  results  lead  to  the  con- 
clusion that  the  atomic  weight  of  an  element  is  not  its  most  charac- 
teristic property.  The  two  forms  of  lead  are  called  isotopes,  the 
name  signifying  that  the  two  varieties  have  different  atomic 
weights,  but  that  in  all  other  properties  they  are  identical.  Since 
the  announcement  of  this  important  discovery  in  the  case  of  lead, 
many  attempts  have  been  made  to  isolate  the  isotopes  of  other 
elements,  but  adequate  experimental  evidence  has  not  been 
furnished  in  any  case.  The  fact  that  isotopes  of  other  elements 
exist,  has  been  demonstrated,  however. 

On  account  of  the  fact  that  the  atomic  weight  of  lead  from 
sources  which  are  not  radioactive  is  207.2,  it  appears  probable 
that  lead  of  this  kind  is  a  mixture  of  isotopes,  one  of  which  is 
derived  from  uranium  and  the  other  of  which  possesses  an  atomic 
weight  greater  than  207.2.  This  conclusion  receives  striking  con- 
firmation from  the  fact  that  radioactive  thorium,  which  has  the 
atomic  weight  232.4,  is  the  first  member  of  a  series  of  radioactive 
elements,  which  also  pass  finally  into  lead;  the  atomic  weight 
of  the  latter  in  this  case  is  208.4,  because  the  changes  involve  the 
loss  of  6  helium  atoms:  232.4  -6X4  =  208.4.  A  determina- 
tion of  the  atomic  weight  of  a  sample  of  lead  obtained  from  a 
thorium  mineral  gave  the  value  208.4.  It  seems  highly  prob- 
able, therefore,  that  ordinary  lead  is  a  mixture  of  two  isotopes, 
one  of  which  is  derived  from  uranium  and  the  other  from 
thorium. 

801.  Radioactive  changes  differ  from  ordinary  chemical 
transformations  in  that  the  rate  at  which  they  take  place  does 
not  vary  with  temperature.  The  chemist  knows  how  to  modify 
the  reactions  that  occur  between  atoms,  but  up  to  the  present  no 
one  has  been  able  to  control  radioactive  disintegrations,  which 
result  from  the  spontaneous  decomposition  of  the  atoms  them- 
selves. 

The  rate  at  which  radium  and  other  radioactive  elements 
decompose  is  proportional  to  the  amount  of  the  substance  under- 
going the  change;  consequently,  as  the  amount  diminishes  the 
rate  gets  slower  and  slower.  In  order  to  express  the  rate  at 
which  the  several  elements  of  this  class  disintegrate  it  is  usual 


RADIOACTIVITY.     THE  STRUCTURE  OF  ATOMS         651 

to  state  the  time  required  for  one-half  of  any  given  amount  to 
decompose.  In  the  case  of  radium  this  time  is  1850  years,  and 
in  the  case  of  niton  3.85  days.  One  atom  of  radium  out  of  each 
100,000,000,000  undergoes  decomposition  per  second. 

802.  There  is  a  marked  difference  in  the  "  life  "  of  other  ele- 
ments of  this  kind;    radium-C  exists  for  a  fraction  of  a  second, 
whereas  in  the  case  of  uranium,  the  period  for  one-half  decom- 
position is   approximately   5  X  109   years.     It  has   been   deter- 
mined that  1  gram  of  uranium  gives  off  helium  at  the  rate  of  1  c.c. 
in  16,000,000  years.     The  mineral  fergusonite  contains  26  c.c. 
of  helium  for  each  gram  of  uranium  present.     If  the  gas  -was 
produced  as  the  result  of  the  disintegration  of  uranium,  the  process 
must  have  been  taking  place  for  416  million  years. 

The  amount  of  energy  set  free  as  heat  when  radioactive  trans- 
formations take  place  is  very  great;  it  has  been  calculated  that 
the  total  energy  liberated  when  an  atomic  weight  of  niton  changes 
to  radium-A  is  equal  to  60,000  horse-power  for  one  day. 

803.  Avogadro's    Number. — The    a-particles    expelled    from 
radium  travel  at  a  velocity  of  about  0.1  that  of  light,  and  when 
they  strike  certain  substances,  zinc  sulphide,  for  example,  they 
render  the  latter  luminescent.     The  instrument  used  to  observe 
this  phenomenon  is  called  a  spinthariscope.     It  consists  of  a  small 
tube  at  one  end  of  which  is  placed  a  trace  of  a  radium  salt  in  front 
of  a  screen  covered  with  zinc  sulphide,  and  at  the  other  end  of 
which  is  a  lens  through  which  the  scintillations  produced  can  be 
seen.     It  is  possible  to  determine  the  number  of  helium  atoms 
given  off  by  radium  in  a  given  time  by  using  a  very  small  amount 
of  the  salt  and  placing  the  screen  at  a  considerable  distance  from 
it.     In  this  way  the  number  of  scintillations  can  be  reduced 
to  such  an  extent  that  they  can  be  readily  counted.     It  was 
determined  that  1  gram  of  radium  gives  off  1.36  X  1011  a-par- 
ticles  per  second,  or  4.29  X  1018  helium  atoms  per  year.     The 
helium  given  off  by  radium  has  been  collected  and  measured, 
and  it  was  found  that  the  gas  was  formed  at  a  rate  of  0.156  c.c. 
per  year.     It  follows,  therefore,  that  0.156  c.c.  of  helium  contains 
4.29  X  1018  atoms,  from  which  can  be  calculated  that  1  c.c.  of 
the  gas  contains  2.75  X  1019  atoms  and  1  gram-atomic-weight 
6.16  X  1023  atoms.     Since  helium  is  a  monatomic  gas,  the  num- 
ber of  atoms  present  in  its  gram-atomic-weight  is  equal  to  the 


652  INORGANIC  CHEMISTRY  FOR  COLLEGES 

number  of  molecules  present  in  the  gram-molecular-weight  of 
any  other  gas;  6.16  X  1023  is,  therefore,  the  number  of  molecules 
in  a  gram-molecular-weight  of  a  gas  (22.4  liters).  This  number 
has  been  arrived  at  in  a  variety  of  independent  ways,  and  there  is, 
therefore,  strong  evidence  of  its  correctness;  it  is  known  as  Avoga- 
dro's  number. 

804.  The  Structure  of  the  Atom. — We  have  seen    that  the 
radioactive   elements   are   constantly    undergoing    disintegration 
as  a  result  of  which  negative  charges  of  electricity  are  produced. 
It  has  also  been  pointed  out  (565)  that  when  a  circuit  made  up 
of  two  metallic  wires  is  heated  at  one  of  the  junctions,  an  electric 
current  is  set  up  as  the  result  of  the  flow  of  electrons  through  the 
wires.     When  a  metallic  wire  is  highly  heated  it  gives  off  electrons 
into  space,  and  if  the  wire  is  placed  in  a  vacuum  tube  which  con- 
tains also  a  sheet  of  metal  that  is  charged  positively,  the  negative 
charges  pass  through  the  vacuum  from  the  wire  to  the  sheet  of 
metal  and  an  electric  current  is  produced.     Such  facts  as  the 
above  lead  to  the  conclusion  that  the  atoms  of  which  matter 
is  composed  are  made  up  of  electrical  charges/  which  become 
evident  when  the  atom  is  heated  or  when  it  spontaneously  decom- 
poses. 

These  views  have  been  given  concrete  expression  by  the 
scientists  who  have  studied  the  phenomenon  outlined  above,  and 
have  led  to  a  conception  of  the  structure  of  atoms  which  is  the 
guiding  principle  in  the  research  of  to-day.  Among  the  inves- 
tigators who  have  carried  out  the  experimental  work  and  have 
formulated  the  theories  are  J.  J.  Thompson,  Rutherford,  Soddy, 
and  Ramsay.  The  discoveries  which  led  to  the  active  investiga- 
tions in  this  field  were  made  by  Crookes  (cathode  rays  in  1878), 
Rontgen  (X-rays,  1895),  Becquerel  (radioactivity  of  uranium 
compounds,  1896),  and  M.  and  Mme.  Curie  (radium,  1903).  A 
large  number  of  physicists  and  chemists  have  aided  in  developing 
the  conception  of  atomic  structure  given  below. 

805.  An  atom  is  thought  to  be  made  up  of  a  central  nucleus 
composed  of  positive  charges  of  electricity  surrounded  by  a  number 
of  negative  charges,  placed  at  a  distance  from  the  nucleus,  which 
is  very  small  compared  with  the  diameter  of  the  space  occupied 
by  the  atom.     The  atoms  differ  from  one  another  in  the  number 
of  positive  and  negative  charges  which  they  contain.     The  weight 


RADIOACTIVITY.     THE  STRUCTURE  OF  ATOMS         653 

of  the  atom  is  associated  with  the  positive  charge.  The  negative 
charges  are  arranged  around  the  nucleus  in  one  or  more  shells,  and 
those  making  up  the  outer  shell  are  capable  of  transference  to 
another  atom;  these  are  the  so-called  valence  electrons,  which 
function  when  chemical  union  between  two  atoms  takes  place. 

This  general  conception  is  applied  to  the  several  elements  as 
follows:  The  hydrogen  atom  is  considered  to  be  made  up  of  one 
positive  charge  and  one  electron,  which  is  in  vibratory  motion 
and  can  function  as  a  valence  electron.  Helium,  the  second 
element,  consists  of  a  positive  nucleus  and  two  electrons  which 
are  so  firmly  bound  that  they  cannot  be  transferred  to  other 
atoms;  the  element  is  inert  chemically.  The  succeeding  elements 
are  made  of  an  inner  core  composed  of  a  positively  charged  nucleus 
and  electrons  which  do  not  function  in  chemical  transformations, 
and  an  outer  sphere  of  valence  electrons,  the  number  of  which 
varies  progressively  as  we  pass  from  one  element  to  the  next  in 
a  series  in  the  periodic  classification.  Lithium  has  1  valence 
electron,  beryllium  2,  boron  3,  carbon  4,  nitrogen  5,  oxygen  6, 
fluorine  7.  Neon,  the  next  element,  has  8,  and  since  it  is  inert, 
the  conclusion  is  drawn  that  this  arrangement  leads  to  such  a 
compensation  of  the  forces  in  the  atom  that  no  electron  can  leave  it. 

In  the  building  up  of  the  next  element,  sodium,  by  the  addition 
of  positive  charges  and  electrons,  those  already  present  and  those 
added,  pass  to  the  inner  core  except  one  electron  which  resides 
on  a  new  outer  shell.  As  we  pass  through  the  series  as  before, 
the  next  element,  magnesium,  has  2  valence  electrons,  aluminium 
3,  and  so  on  to  chlorine,  which  has  7,  and  to  argon  with  8,  which, 
like  neon,  with  8  electrons  in  the  outer  shell,  is  inactive.  The 
building  up  is  continued  through  the  elements  in  this  way,  and  as 
a  result  there  is  a  recurrence  of  elements  with  from  1  to  7  valence 
electrons.  The  elements  thus  exhibit  a  periodicity  in  the  number 
of  valence  electrons  they  possess.  The  position  of  the  electrons  in 
the  outer  sphere  is  indicated  on  page  654.  In  each  case,  however, 
the  inner  core  of  the  atom,  which  is  not  represented,  is  different. 

According  to  the  electronic  conception  of  the  atom,  when  two 
elements  combine,  a  valence  electron  passes  from  one  element  to 
the  other.  When,,  for  example,  sodium  chloride  is  formed,  the 
metal  loses  its  electron,  which  takes  up  its  place  in  the  outer  sphere 
of  the  chlorine  atom;  the  two  atoms  are  held  together  by  the 


654 


INORGANIC  CHEMISTRY  FOR  COLLEGES 


electrical  field  of  force  set  up  between  them  as  the  result  of  the 
transfer.  When  the  resulting  compound  is  dissolved  in  water  and 
the  atoms  are  separated,  the  sodium  having  lost  an  electron,  which 
is  a  negative  charge,  becomes  a  positively  charged  ion;  and  the 
chlorine  having  received  an  additional  negative  charge  becomes 
a  negative  ion. 


FIG.  44. 

806.  Combination  between  atoms  can  take  place  in  a  different 
way  from  that  just  described.  The  two  atoms  may  share  elec- 
trons so  that  neither  loses  any.  When  2  fluorine  atoms  unite  to 
form  a  molecule,  F2,  it  is  assumed,  according  to  this  theory,  that 
the  atoms  so  arrange  themselves  that  they  share  2  electrons  in 
common  in  such  a  way  that  two  groups  of  8  electrons  are  formed. 
This  is  represented  in  Fig.  45.  Each  fluorine  atom  has  7  elec- 
trons, and  the  arrangement  of  the  two 
atoms  indicated  has  14  electrons;  two 
electrons  are  shared  in  common  by  the  two 
atoms.  Compounds  of  this  kind  do  not 
have  the  field  of  force  between  the  two 
atoms  that  is  set  up  in  the  way  indicated 
in  the  case  of  sodium  chloride;  they  are 
not  ionized  and  are  very  stable.  There 

appears  to  be  a  tendency  for  atoms  to  arrange  themselves  so  that 
8  electrons  are  present  in  the  outer  sphere;  and,  as  we  have  seen, 
this  configuration  is  the  most  stable  and  is  present  in  the  inert 


FIG.  45. 


RADIOACTIVITY.     THE  STRUCTURE  OF  ATOMS          655 

The  positive  valence  that  an  atom  can  show  is  determined  by 
the  number  of  electrons  it  has  in  its  outer  sphere,  because  when 
an  element  functions  as  a  positive  element  it  loses  electrons  and, 
therefore,  assumes  a  positive  charge.  The  negative  valence  is 
determined  by  the  difference  between  8  and  the  number  of  valence 
electrons,  for  the  maximum  number  of  the  latter  that  can  be 
present  in  the  outer  sphere  is  8.  If,  for  example,  an  element 
possesses  7,  as  in  the  case  of  chlorine,  it  can  take  up  but  1  electron; 
its  negative  valence  is,  accordingly  1.  On  account  of  the  signifi- 
cance of  the  number  eight  this  theory  is  called  the  octet  theory  of 
valence.  It  has  been  developed  by  G.  N.  Lewis  and  also  by 
Langmuir,  who  has  used  it  to  correlate  many  chemical  facts. 

807.  It  will  be  seen  from  the  above  that  in  ordinary  chemical 
phenomena  the  outer  sphere  of  the  atom  is  alone  concerned. 
The  inner  core  maintains  its  identity  in  all  such  transformations, 
and  the  elements  as  we  know  them  pass  from  one  compound  to 
another  without  change.  In  radioactive  transformations,  the 
inner  core  of  the  atom  undergoes  change  and,  as  a  result  new  ele- 
ments are  formed.  Since  helium  is  given  off  in  many  of  these 
changes,  we  must  conclude  that  it  is  a  constituent  of  all  these 
atoms,  and  the  reasonable  assumption  is  that  all  atoms,  except 
hydrogen,  contain  helium.  Since  the  atom  of  helium  weighs  4 
and  the  atomic  weights  of  the  elements  are  not  multiples  of  this 
number,  we  must  conclude,  further,  that  a  lighter  atom  must  be 
present  in  the  complex  atom  of  many  elements.  The  most  rea- 
sonable view  to  take  is  that  hydrogen  is  the  other  element  involved. 
Inspired  by  this  hypothesis  Rutherford  recently  subjected  pure 
nitrogen  to  a  bombardment  by  a-particles  and  obtained  as  a 
result  a  trace  of  hydrogen.  No  hydrogen  was  obtained  from  oxy- 
gen. These  results  are  significant  in  view  of  the  fact  that  the 
atomic  weight  of  oxygen,  16,  is  a  multiple  of  4,  whereas  that  of 
nitrogen,  14,  could  not  be  made  up  solely  of  helium. 

It  is  worthy  of  note  that  the  elements  which  are  sponta- 
neously undergoing  radioactive  disintegration  are  those  of  high 
atomic  weight — the  ones  that  contain  a  large  number  of  the  simpler 
units  of  which  the  elements  are  composed.  It  is  possible,  there- 
fore, that  the  decomposition  of  the  heavier  atoms  that  do  not 
show  radioactivity  may  be  effected  more  readily  than  that  of  the 
light  elements,  when  the  means  to  effect  such  changes  are  at  hand. 


656  INORGANIC  CHEMISTRY  FOR  COLLEGES 

808.  General  Conclusions. — From  the  very  brief  and  ele- 
mentary account  of  the  investigations  of  the  phenomena  of  radio- 
activity which  has  been  given,  it  will  be  seen  that  it  is  now  pos- 
sible to  study  the  complex  nature  of  the  atoms,  which  up  to  recent 
years  have  been  the  final  chemical  units  of  matter.  The  con- 
ception of  these  units  has  changed  as  matter  has  been  studied, 
and  as  the  means  by  which  matter  can  be  decomposed  have  been 
learned.  When  heat  energy  alone  was  available,  substances  that 
are  not  decomposed  by  heat  at  the  temperature  obtained  by  chem- 
ical reactions,  were  thought  to  be  elementary.  When  electrical 
energy  was  discovered  and  it  was  learned  how  to  make  and  apply 
it,  many  substances  undecomposed  before  were  shown  to  be  com- 
plex. Lime,  for  example,  yielded  a  new  metal — calcium.  And 
now,  the  energy  developed  when  radioactive  transformations 
take  place  carries  us  one  step  farther;  when  it  is  applied  to  the 
atom  that  resists  ordinary  electrical  energy,  the  atom  yields  and 
decomposition  takes  place. 

Electricity  was  known  a  long  time  before  it  was  found  out  how 
to  generate  it  from  other  kinds  of  energy  and  how  to  use  it  to 
effect  the  decomposition  of  chemical  compounds.  It  is  not 
unreasonable  to  think  that  a  way  will  be  found  to  control  or  per- 
haps produce  radioactivity,  and  when  this  happens  practical 
transmutation  of  the  elements  may  be  possible. 

The  elements  are  a  storehouse  of  uncalculable  energy,  and  if 
a  way  can  be  found  to  bring  about  their  decomposition  this  energy 
can  be  made  available  for  the  uses  of  man. 

The  study  of  radioactivity  has  led  not  only  to  the  view  that 
atoms  are  complex,  but  has  shown  that  one  of  the  fundamental 
laws  upon  which  the  quantitative  study  of  chemical  phenomena 
is  based,  is  not  an  exact  statement  of  the  truth.  The  law  of  def- 
inite proportions  does  not  apply  to  isotopes;  the  compounds 
made  from  lead  obtained  from  radioactive  sources  do  not  contain 
a  fixed  proportion  of  the  elements  present.  In  one  case  208.4 
grams  of  lead  unite  with  2X35.45  grams  of  chlorine  to  form  lead 
chloride  and  in  the  other  case  206.4  grams  of  lead  unite  with  the 
same  weight  of  the  halogen.  But  the  newer  knowledge  does  not  in 
any  way  modify  the  facts  which  have  been  learned  as  the  result 
of  the  intensive  study-  of  chemical  phenomena;  they  teach  us 
that  there  is  much  ahead  to  be  discovered  and  that  we  must  mod- 


RADIOACTIVITY.     THE  STRUCTURE  OF  ATOMS         657 

ify  our  conclusions  and  theories  as  the  science  grows.  The  greater 
the  modifications  necessary,  the  greater  the  growth  of  the  science. 
The  last  decade  has  advanced  our  knowledge  of  matter  to  such  a 
remarkable  degree  that  it  is  impossible  to  foresee  what  lies  just 
ahead. 


APPENDIX 


I.  TEMPERATURE  SCALES 

On  the  centigrade  scale,  the  freezing-point  and  boiling-point  of  water 
are  called,  respectively,  0°  and  100°.  On  the  Fahrenheit  scale,  these  points 
are  called,  respectively,  32°  and  212°.  Therefore,  100  degrees  C.  =  212  —  32 
=  180  degrees  F.  One  degree  C.  =  f  degree  F.  and  1  degree  F.  =  f  degree  C. 
Temperatures  on  one  scale  can  be  changed  to  temperatures  on  the  other  by 
the  use  of  the  following  formulas,  in  which  C°  is  the  reading  on  the  centigrade 
thermometer  and  F°  is  the  reading  on  the  Fahrenheit  thermometer. 

C°  =  f  (F°  -  32),         F°  =  f  (C°)  +  32 
Absolute  temperature,  T°,  =  C°+  273°. 

II.  METRIC  SYSTEM 

Length. — 1000  Millimeters  =  100  centimeters  =  1  meter  =  39. 37 inches. 
1  Inch  =  2.54  centimeters. 

Volume. — 1000  Cubic  centimeters  =  1  liter  =  0.03532  cu.  ft.  =  61.03 
cu.  in.  =  1.057  quarts  =  34.1  fluid  ounces. 

1  Fluid  ounce  =  29.57  c.c.  1  Quart  =  0.9434  liter.  1  Cubic  foot  =  28.32 
liters. 

Weight. — 1000  milligrams  =  100  centigrams  =  10  decigrams  =  1  gram 
=  0.03527  ounce  =  15.43  grains. 

1  Kilogram  =  1000  grams  =  2.205  pounds  avoirdupois. 

1  Pound  avoirdupois  =  453.6  grams. 

1  Ounce  avoirdupois  =  28.35  grams. 

1  Metric  ton  =  1000  kilos  =  2205  Ibs. 

III.  ENERGY  UNITS 

1  Calorie  (cal.)  =  heat  to  raise  the  temperature  of  1  gram  of  water  1 
degree  centigrade  (at  15°).  1000  Calories  =  1  large  calorie  (Cal.). 

1  British  thermal  unit  (B.t.u.)  =  heat  required  to  raise  the  temperature 
of  1  Ib.  of  water  1  degree  Fahrenheit  at  its  maximum  density  =  252  cal. 

659 


660 


APPENDIX 


Joules  =  coulombs  X  volts.  1  Joule  =  0.238  cal.  1  Calorie  =  4.187 
joules.  1  B.t.u.  =  1055  joules. 

IWatt  =  1  joule  per  second.  1  Horse-power  =  33,000  ft.-lbs.  per 
minute  =  746  watts. 


IV.  MOLECULAR  VOLUMES 

1  Gram-molecular-weight  of  a  gas  occupies  22.4  liters  at  0°  and  760  mm. 
1  Pound-molecular-weight  of  a  gas  occupies  359  cubic  feet  at  0°   and 
760  mm. 


V.  VAPOR  PRESSURE  OF  WATER 

The  temperatures  are  given  in  degrees  centigrade,  and  the  pressure  in 
millimeters  of  mercury. 


Temp. 

Pressure. 

Temp. 

Pressure. 

Temp. 

Pressure. 

0° 

4.6 

16° 

13.5 

26° 

25.1 

5 

6.5 

17 

14.4 

27 

26.5 

8 

8.0 

18 

15.4 

28 

28.1 

9 

8.6 

19 

16.3 

29 

29.8 

10 

9.2 

20 

17.4 

30 

31.5 

11 

9.8 

21 

18.5 

31 

33.4 

12 

10.5 

22 

19.7 

32 

35.4 

13 

11.2 

23 

20.9 

33 

37.4 

14 

11.9 

24 

22.2 

34 

39.6 

15 

12.7 

25 

23.6 

35 

41.8 

VI.  GROUPS  OP  THE  COMMON  METALS  IN  QUALITATIVE  ANALYSIS 

Group  1.     Precipitated  by  HC1  :  AgCl,  HgCl,  PbCl2. 

Group  2.  Precipitated  by  H2S  in  HC1  solution:  HgS,  CuS,  PbS,  Bi2S3. 
AsjSa,  Sb2S3,  SnS.  The  last  three  dissolve  in  (NH4)2Sz. 

Group  3.  Precipitated  by  (NH4)2S  from  neutral  solution:  FeS,  CoS, 
NiS,  MnS,  ZnS,  FeS  +  S  from  ferric  salts,  Cr(OH)3,  A1(OH)3.  Trivalent 
metals  may  be  separated  before  addition  of  (NH4)2S  by  precipitating  as 
hydroxides  with  NH4OH  in  presence  of  NH4C1.  Addition  of  (NH4)2S  to 
filtrate  precipitates  sulphides  of  bivalent  metals. 

Group  4.  Precipitated  by  (NH4)2COa  :  BaCO3,  SrCO3,  CaCO3.  Ammo- 
nium phosphate  precipitates  from  filtrate  MgNH4PO4.  Solution  contains 
K,  Na,  and  NH4  salts. 


APPENDIX  661 


VII.  ELECTROMOTIVE  SERIES  OF  THE  ELEMENTS 

Potassium  Iron  Silver 

Sodium  Cobalt  Palladium 

Barium  Nickel  Platinum 

Strontium  Tin  Gold 

Calcium  Lead  

Magnesium  Hydrogen  Iodine 

Aluminium  Copper  Bromine 

Manganese  Arsenic  Oxygen  (O       ) 

Zinc  Bismuth  Chlorine 

Chromium  Antimony  Oxygen  (OH  ~) 

Cadmium  Mercury 


INDEX 


(The  references  are  to  pages;  the  more  important  ones  are  given  in  bold- 
face type.) 


Acetic  acid,  425 
Acetylene,  34,  192,  423 
ACHESON,  174 
Acids,  41,  238,  425 

action  on  dyes,  331 
metals,  458,  460 
salts,  507 

organic,  425 
Air,  analysis,  290 

content  of  water-vapor  in,  91 

liquid,  166 

latent  heat,  168 
ALBERTUS  MAGNUS,  5 
Alchemy,  4 
Alcohol,  422,  424 

heat  of  vaporization,  162 

wood,  425 
Alkalies,  210 

action  on  metals,  465 
ALLEN-MOORE  cell,  519 
Alloys,  449 

composition,  455 

fusible,  417 
Allo tropic  forms  of  carbon,  178 

of  elements,  147 
Alum,  571 
Aluminium,  457,  466,  666 

and  acids,  459,  460 

bronzes,  568 

chloride,  570 

hydroxide,  569 

oxide,  569 

sulphate,  570 

hydrolysis,  513,  514 
.      tests  for,  570 
Alundum,  569 
Alunite,  527 


Amatol,  333 
Ammonia,  302 

absorption  by  charcoal,  176 

action  on  salts,  509 

chemical  properties,  311 

cyanamide  process,  309 

formation,  310 

heat  of  formation,  306 

heat  of  vaporization,  162 

history,  303 

in  coal  gas,  200 

in  water,  336 

occurrence,  303 

oxidation,  313,  325 

physical  properties,  311 

preparation,  303 

synthetic,  304 

uses,  314 
Ammonium  bicarbonate,  316 

carbonate,  316 

chloride,  315 

dichromate,  630 

hydroxide,  313,  502 

molybdate,  636 

nitrate,  333 

nitrite,  287 

phosphomolybdate,  404,  636 

polysulphide,  254 

radical,  315 

salts,  314 

tests  for,  316 

sulphate,  316 

solubility,  506 

thioarsenate,  412 

thioarsenite,  412 

thiocyanate,  618 
AMPERE,  473 
Amphoteric  elements,  413,  440 

663 


664 


INDEX 


Aniline,  330 

Antidotes  for  poisons,  409 
Antimonic  acid,  415 
Antimonous  acid,  415 
Antimony,  413,  414 

chlorides,  415 

halides,  416 

oxides,  415 

salts,  415 

sulphides,  416 
Aqua  regia,  329 
Aqueous  vapor,  85 
Aragonite,  533 
Argon,  295,  301 

use,  298 
ARRHENIUS,  215 
Arsenic,  407 

acids  of,  409 

halides,  411 

pentoxide,  409 

properties,  408 

sulphides,  411 

trioxide,  408 

trisulphide,  412 
Arsine,  410 
Arsenite,  407 
Asbestos,  432 

Association  of  molecules,  379 
Atmosphere,  284,  289 
Atom,  structure  of,  652 
Atomic  numbers,  646 

theory,  52 

weights,  determination,  345,  348, 

352 

Atoms,  number  in  gram-molecule,  156 
Auric.     See  also  gold. 

chloride,  heat  of  formation,  143 
AVOGADRO,  law  of,  346 
AVOGADRO'S  number,  651 
Azurite,  594 


Babbit  metal,  454,  455 
BACON,  ROGER,  5 
Bacteria,  114,  295 
nitrifying,  288 


Baking  powders,  524 
BALARD,  365 
Barium,  531,  646 

chloride,  547 

hydroxide,  546 
solubility,  531 

oxides,  546 

peroxide,  150,  152 

salts,  547 

sulphate,  547 

solubility,  531 

tests  for,  548 
Barometer,  79 
Bases,  210 

Batteries,  storage,  490,  493 
BECQUEREL,  548 
Bell  metal,  455 
Benzene,  199 
Benzoic  acid,  426 
Beryllium,  551 
BERZELIUS,  392 
BESSEMER  process,  608 
BIRKELAND-EYDE  process,  324 
Bismuth,  417 

glance,  417 

ocher,  417 

oxides,  418 

salts,  418 

sulphide,  418 

tests  for,  419 
BLACK,  179 
Blast  furnace,  604 
Bleaching,  113,  153,  148,  262,  280 

powder,  104,  542 
Blow-pipe,  oxy-hydrogen,  49 
Boiler  scale,  554 
Boiling-point,  160 

of  solutions,  224 
Bone-black,  177 
Boracite,  435 
Borates,  436 
Borax,  435,  436 

-beads,  438,  618 

hydrolysis,  514 
Bordeaux  mixture,  598 
Boric  acid,  434,  435 

tests  for,  439 


INDEX 


665 


Bornite,  594 
Boron,  434,  435 

carbide,  435 

hydrides,  439 

nitride,  435 
BOYLE,  80 

law  of,  285 
BRAND,  397 
Brass,  453,  455 
Braunite,  639 
British  thermal  unit,  139 
Britannia  metal,  455 
Bromic  acid,  387 

Bromides,  properties  and  test  for,  369 
Bromine,  365 

chemical  properties,  367 

hydrate,  367 

occurrence,  365 

physical  properties,  367 

preparation,  366 

uses,  367 
Bronze,  455 

Bronzes,  aluminium,  568 
Butyric  acid,  426 
By-product  ovens,  199 


Cadmium,  558 

salts,  558 
Caesium,  515 
Calamine,  555 
Calcite,  533 
Calcium,  531 

bicarbonate,  534 

carbide,  191 

carbonate,  533 

chloride,  99,  532 
hydrated,  94 

cyanamide,  310 

hydroxide,  538 
solubility,  531 

-light,  50 

oxide,  537 

heat  of  formation,  142 

phosphate,    solubility   in   acids, 
507 

phosphates,  540 


Calcium,  phosphide,  399 
sulphate,  539,  554 
solubility,  531 
sulphide,  543 
tests  for,  545 

Calculations,  chemical,  68 
Calorie,  90,  139,  159 
Calorimeter,  139 
Carbides,  179 
Carbohydrates,  423 
Carbolic  acid,  330 
Carbon,  171 

allotropic  forms,  178 
burning  of,  temperature,  142 
chemical  properties,  179 
cycle  of,  186 
physical  properties,  179 
dioxide,  179,  182 

critical  temperature,  165 
heat  of  formation,  141 
history,  179 
in  the  air,  291 
in  nature,  186 
occurrence,  179 
physical  properties,  179 
preparation,  180 
test  for,  184 
uses,  185 
disulphide,  190 
monoxide,  187 

chemical     properties,     177, 

189 

physical  properties,  189 
preparation,  188 
test  for,  190 
prints,  630 

tetrachloride,  115,  191 
Carbona,  191 
Carbonates,  preparation,  properties 

and  test  for,  184 
Carborundum,    194 
Carnalite,  526 
CARO'S  acid,  281 
CASTNER-KELLNER  cell,  519 
Catalysis,  26,  101,  177,  307 
CAVENDISH,  38,  88,  297,  323 
Cell,  Daniell,  479 


INDEX 


Cell,  dry,  480 

electric,  479 

gravity,  480 

Leclanche",  481 
Cellulose,  423,  596 

nitrate,  427 
Cement,  575 
Cementite,  610 
Chalcocite,  594 
Chalcopyrite,  594 
Chalk,  533 
Charcoal,  animal,  177 

heat  of  combustion,  140 

wood,  175,  176 
CHARLES,  law  of,  82 
Chemical  equilibrium,  233 

reactions,  types  of,  130 
Chlorates,  384,  385 
Chloric  acid,  384,  385 
Chlorides,  behavior  when  heated,  115 

formation,  128 

test  for,  128,  130 
Chlorine,  99 

action  with  water,  110 

as  bleaching  agent,  113 

chemical  properties,  106 

comparison  of  with  oxygen,  115 

critical  temperature,  165 

dioxide,  386 

history,  100 

hydrate,  109 

liquefaction,  109 

occurrence,  99 

oxidizing  action,  111 

physical  properties,  105 

preparation,  100-105 

test,  114 

uses,  114 

Chlorine  monoxide,  382 
Chloroform,  130 
Chlorophyl,  26,  187 
Chrome  alum,  593,  628 
Chromates,  628 
Chromel,  494 
Chromic  anhydride,  629 

chloride,  628 

compounds,  627 


Chromic  oxide,  628 
Chromium,  626 

in  steel,  612 

test  for,  635 
Chromite,  626 
Chromous  compounds,  627 
Citric  acid,  426 
CLAUDE,  307 
Clay,  571 

Climate,  effect  of  water-vapor  on,  89 
Coal,  194 

analysis,  195 

anthracite,  196 

bituminous,  196 

burning  of,  196 

formation,  194 

semi-bituminous,  196 

yield  of  gas  from,  202 
Coal  tar,  200 
Cobalt,  619 

tests  for,  620 
Cobaltic  compounds,  619 
Cobaltite,  408 
Cobaltous  compounds,  619 
Coins,  601 
Coke,  198 

analysis  of,  196 
Colemanite,  435 
Colloidal  clay,  574 

metals,  447 
Combination,  131 
Combustion,  31 

heat  of,  140 

long-flame,  207 

spontaneous,  33 

surface,  206 

Compounds,  chemical,  17 
Conduction,  electrolytic,  489 
Conductivity,  electrical  of  metals,  446 

heat,  446 

Copper,  463,  594.     See  also  cuprous 
and  cupric 

action  of  acids  on,  462 

ferro cyanide,  614 

hydroxide,  596 

metallurgy,  469 

nitrate,  341,  598 


INDEX 


667 


Copper  oxides,  595,  596 

purification,  495 

reaction  with  nitric  acid,  327 

sulphate,  93,  94,  598,  513 

sulphide,  599 

tests  for,  599 
Coral,  533 
Corrosion  of  iron,  617 

of  metals,  462 
Corundum,  569 
Cotton,  596 
COULOMB,  473 
Coulometer,  483 
Couples,  metallic,  458,  478 
COURTOIS,  370 
Critical  temperature,  164 
Crocoisite,  627 
CROOKES,  652- 
Croton  water,  analysis,  337 
Cryolite,  570 
Crystallography,  157 
Cupric  bromide,  597 

chloride,  597 
Cuprite,  594 
Cuprous  chloride,  597 

cyanide,  598 

iodide,  597 
CURIE,  548,  652 

D 

DALTON,  54-56,  346,  358 

DAVY,  100,  338,  527 

DEACON  process,  101 

Decomposition,  double,  131,  134,  221 

Deliquescence,  126 

Density,  159,  442 

Detonators,  331 

Dew,  90 

DEWAR,  109 

Diamond,  140,  172,  178 

Dichromates,  629 

Diffusion,  155 

of  gases,  rate,  45 
Disilicic  acids,  431 
Distillation,  96,  163 
Dissociation,  113 

electrolytic,  215 


Double  decomposition,  131,  134,  221 
DULONG,  352 
DUMAS,  109 
Duriron,  464,  612 
Dust  hi  the  air,  295 
Dyes,  bleaching,  113 
sulphur,  249 

E 

EDISON  storage  cell,  493 
Efflorescence,  93 
ELDRED,  207 
Electricity,  8 
Electrochemistry,  471 
Electrolysis,  217 

of  water,  27 
Electrolytes,  216 
Electrolytic  conduction,  489 
Electromotive  force  of  cells,  488 

series  of  elements,  223,  486,  488 
Electronic  valence,  389 
Electrons,  389,  472,  482,  647 
Electroplating,  464,  496 
Electrotypes,  496 
Elements,  19 

acid-forming,  439 

amphoteric,  413,  440 

metallic,  209 

non-metallic,  210 
Emery,  569 
Emulsions,  437 
Enamels,  437,  573 
Endothermic  reactions,  141 
Energy,  11,  137 

electrical,  472 

free,  137 
Epsom  salt,  551 

Equations,  writing  of,  61,  103,  132, 
257-259 

partial,  258 

thermo-chemical,  141 
Equlibrium,  chemical,  233 
Esters,  426 
Ethane,  421 
Ethyl  acetate,  426 

borate,  439 

chloride,  338 


668 


INDEX 


Ethylene,  422 
Eutectic,  451 
Exothermic  reactions,  141 
Explosions,  47,  330,  342 


Family,  chemical,  413,  439,  441 
FAHRENHEIT,  78 
FARADAY,  109,  164,  165 

law  of,  482 
Fats,  426 

cooking,  50 

heat  of  combustion,  140 
FEHLING'S  solution,  596 
Feldspar,  432 
Fermentation,  424,  425 
Ferric  acetate,  618 

-ammonium  alum,  511 

chloride,  614 

compounds,  614 

hydroxide,  503,  615 

oxalate,  593 

oxide,  615 

sulphate,  616 
Ferrite,  611 
Ferrous-ammonium  sulphate,  613 

carbonate,  613 

chloride,  613 

compounds,  613,  614 

sulphate,  613 

sulphide,  614 
Fertilizers,  541 
Fire  extinguishers,  115,  185 
Flame,  Bunsen,  temperature,  204 

hydrogen,  temperature,  142 
Flames,  205 
Flint,  430 
Flotation,  595 
Fluorides,  381 
Fluorine,  375 

discovery,  377 

occurrence,  376 

preparation,  377 

properties,  378 
Flux,  437,  468 
Formalin,  425 
Formulas,  58 


Formulas  and  valence,  120 

determination,  349 

graphic,  119,  389 
FOWLER'S  solution,  409 
Franklinite,  555 
FRASCH,  246 
Freezing  mixtures,  505 

-point  of  solutions,  227 
Fructose,  424 
Fuels,  analyses,  196 

heat  value,  196 
Fuller's  earth,  574 
Furnaces,  electric,  494 

G 

Galena,  582 
Garnet,  431 
Gas-black,  178 

burning,  of,  204 

coal,  200 

illuminating,  200 

calorific  value,  201 

masks,  643 

natural,  300,  422 

producer,  202 

water,  201 
Gases,  155 

adsorption,  176 

calculation  of  volumes,  72 

kinetic  theory,  156 

liquefaction,  164 

measurement,  77 

molecular  volume,  75 

reduction  of  volume,  84,  86 

relative  densities,  75 

solubility,  166 
Gasoline,  115,  422 
GAY-LUSSAC,  322,  370,  377 

law  of,  345 
Germanium,  578 
GLADSTONE,  239 
Glass,  409,  543 

history,  3 

and  hydrofluoric  acid,  381 
GLAUBER,  123 
GLAUBER'S  salt,  124 
Glucose,  424,  596 


INDEX 


669 


Glycerine,  426 
Gold,  599 

tests  for,  602 

Gram-molecular-volume,  74 
Gram-molecular-weight,  59,  74 
Granite,  432 
Graphite,  artificial,  173,  174,  192 

heat  of  combustion,  178 
Gravity,  specific,  159 
Guano,  321 
GULDBERG,  239 
Gun  metal,  455 

powder,  528 
Gypsum,  539 

H 

HABER,  305,  308 
HALL,  567 
Halogen  family,  364 

and  periodic  law,  388 
Hardness,  442 
Hausmannite,  639 
Heat,  7 

energy,  measurement,  138 

of  formation,  141 
of  chlorides,  143 
of  oxides,  142 

of  vaporization,  161 

specific,  159 

Helium,  299,  301,  561,  653 
HELMHOLTZ,  137 
Hematite,  602 
Hemoglobin,  603 
HENRY,  law  of,  166 
VAN'T  HOFF,  law  of,  242 
HOFMANN  apparatus,  9 
Honey,  424 
Horse-power,  473 
Hydrates,  92,  109 
Hydra  zine,  317 
Hydrazoic  acid,  318 
Hydriodic  acid,  373,  374 
Hydrobromic  acid,  307,  368 
Hydrocarbons,  action  with  chlorine, 

108 
Hydrochloric  acid,  123 

chemical  properties,  125-127 


Hydrochloric  acid,  commercial,  127 

constant  boiling  solution,  128 

electrolysis,  100 

physical  properties,  124,  127 

preparation,  124 
Hydrocyanic  acid,  529 
Hydrofluoric  acid,  378,  380 
Hydrofluosilicic  acid,  434 
Hydrogen,  38 

action  with  chlorine,  107 

bromide.    See  hydrobromic  acid. 

chemical  properties,  46,  48 

chloride.     See  hydrochloric  acid. 

critical  temperature,  165 

explosion,  47 

history,  38 

in  the  air,  294 

iodide.     See  hydriodic  acid. 

molecular  weight  of,  156 

occurrence,  38 

peroxide,  149-153,  294 

physical  properties,  44-46 

preparation,  39-44 

selenide,  393 

sulphide,  249,  251 
test  for,  254 

telluride,  393 

temperature  of  flame,  49 

uses,  49 
Hydrolysis,  252 

of  salts,  511 
Hydroxides,  501 
Hydroxylamine,  318 
Hypo,  593 

-bromous  acid,  387 

-chlorites,  383 

-chlorous  acid,  382,  383 

-iodous  acid,  387 

-phosphorous  acid,  405 
Hypothesis,  57 

I 

latrochemists,  413 
Ice,  heat  of  fusion,  167 
lodic  acid,  387 

anhydride,  387 
Iodides,  test  for,  375 


670 


INDEX 


Iodine,  370-372 

chlorides,  387 

pentoxide,  387 
lodoform,  373 
lonization,  extent  of,  228,  231 

of  salts,  511 

Ion,  common,  effect  of,  509 
Ions,  216 

Iron,  447,  603.     See  also  ferrous  and 
ferric. 

action  of  acids  on,  460 

as    acid-forming     element,    616 

cast,  606 

corrosion  of,  463,  617 

metallurgy,  603 

passive,  460 

properties,  612 

purification,  495 

tests  for,  618 

wrought,  607 
Isomorphism,  503 
Isotopes,  650,  656 


JOULE,  473 


K 


Kainite,  526 
Kaolin,  432 
Kerosene,  422 
Kinetic  theory,  156,  160 
Krypton,  301 


Lactic  acid,  425 

Lactose,  424 

Lampblack,  178 

LANDOLT,  53 

LAUE,  647 

LANGMUIR,  655 

Lanolin,  528 

LAVOISIER,  25,  52,  256,  284,  322 

Law,  57 

of  Avogadro,  346 

of  Boyle,  80 

of  Charles,  82 

of  conservation  of  energy,  137 


Law,  of  conservation  of  matter,  53 

of  definite  proportions,  54 

of  Dulong  and  Petit,  352 

of  Faraday,  482 

of  Gay-Lussac,  345 

of  Henry,  166 

of  Le  Chatelier,  243 

of  molecular  concentration,  238 

of  multiple  proportions,  54 

of  Stephan,  400 

of  van't  Hoff,  242 

periodic,  358 
Lead,  582 

and  acids,  459 

atomic  weight,  649 

carbonate,  585 

chloride,  584 

chromate,  627,  629 

hydroxide,  584 

iodide,  584 

oxides,  583 

salts,  586 

storage  battery,  491 

tests  for,  586 
Leather,  571 
LEBLANC,  520,  521 
LE  CHATELIER,  243,  306 
LEWIS,  G.  N.,  655 
LIEBIG,  267 
Light,  7 
Lime,  207,  537 

chloride  of,  104 

-light,  50 

slaked,  538 
Limestone,  533 
Limonite,  603 
Liquefaction  of  gases,  164 
Liquids,  156 
Lithium,  515 
Lithopone,  547 
LOCKYER,  299 
Lodestone,  615 
LULLUS,  RAYMOND,  5 

M 

Magnesia  alba,  551 
Magnesium,  457,  551 


INDEX 


671 


Magnesium-ammonium      phosphate, 
404,  555 

burning  of,  400 

carbonate,  554 

chloride,  99,  553 

hydroxide,  502,  553 

oxide,  552 

pyrophosphate,  555 

sulphate,  553,  554 

tests  for,  555 
Magnetite,  603 
Malachite,  594,  598 
Manganates,  642 
Manganese,  639 

dioxide,  101 

in  steel,  612 

tests,  645 

Manganic  compounds,  641 
Manganite,  639 
Manganites,  641 
Manganous  compounds,  640 
Marble,  533 
MARSH'S  test,  410 
Mass  law,  239 
Matter,  11 
MAYER,  137 
MAYOW,  284 
MEKER  burner,  204 
Melting-point,  167 
MENDELEJEFP,  359 
Mercuric  chloride,  heat  of  formation, 
143 

oxide,  heat  of  formation,  142 
Mercury,  559 

cyanide,  564 

fulminate,  564 

halides,  561 

nitrates,  560 

oxides,  560 

sulphides,  563 

symbol  for,  65 

tests  for,  564 
Meta-acids,  264 
Metallic  couples,  458,  478 
Metallurgy,  457,  467 
Metals,  action  with  acids,  42,  43,  222, 
458 


Metals,  action  on  alkalies,  43,  464 

action  on  salts,  222,  465 

boiling-   and   melting-points   of, 
448 

chemical  properties,  457 

colloidal,  447 

corrosion,  462 

electromotive  series,  486 

occurrence,  466 

passive,  460 

physical  properties,  442 

physical  state,  446 
Metastannic  acid,  581 
Metathesis,  131 
Methane,  108,  421,  422 
Methyl  alcohol,  425 

ether,  422 
Mica,  431 

Microcosmic  salt,  438 
Mill  scale,  618 
Mispickel,  408 
MOISSAN,  172,  179,  377 
Mol,  definition  of,  59 
Molar  solutions,  228 
Molecular  weight,  59 
Molybdenite,  636 
Molybdenum,  636 
Molybdic  anhydride,  636 
Monel  metal,  455 
MORLEY,  356 
Mortar,  538 
MOSELEY,  647 
Muriatic  acid,  100 

N 

Nascent  state,  112 
NELSON  cell,  519 
Neon,  301 
Neutralization,  219 
Nichrome,  494 
Nickel,  465,  620 

catalysis  with,  50 

compounds  of,  620,  621 

in  steel,  612 

tests  for,  621 
Nitrates,  331 

in  water,  336 


672 


INDEX 


Nitrates,  occurrence,  321 

test  for,  332 
Nitric  acid,  321 

action  on  skin,  331 

action  with  metals,  327,  460 

chemical  properties,  326 

history,  321,  322 

in  the  air,  294 

physical  properties,  326 

preparation,  322 

synthesis,  323 

test  for,  332 
Nitric  anhydride,  343 

oxide,  339 

heat  of  formation,  323 
preparation,  340 
properties,  340 
Nitrites,  occurrence,  336 

test  for,  335 
Nitro-benzene,  330 

-cellulose,  330 

-glycerine,  330 

-toluene,  330 
Nitrogen,  284 

chemical  properties,  288 

history,  285 

occurrence,  286 

physical  properties,  288 

preparation,  287 

uses,  289 
Nitrogen  chloride,  387 

dioxide,  341 

dioxide,      chemical     properties, 
342 

iodide,  387 

oxides,  337 

pentoxide,  343 

tetroxide,  341 
Nitrosyl  chloride,  329,  341 
Nitrous  acid,  333,  335 

anhydride,  334,  342 

oxide,  337-339 

Non-metals,      electromotive     series, 
488 

reactions  with  salts,  223 
Normal  solutions,  229 


O 

Ocean,  salts  in,  96 
Oleic  acid,  426 
Olein,  426 
ONNES,  300 
Opal,  430 
Orpiment,  407,  411 
Orthoclase,  432 
Orthophosphates,  402-404 
Orthophosphoric  acid,  402 
Osmium,  625 
Ovens,  bee-hive,  199 

by-product,  199 
Oxalic  acid,  426 
Oxidation,  49,  631,  632,  644 

slow,  32 
Oxides,  499 
Oxygen,  21 

chemical  properties,  30 

comparison  of  with  chlorine,  15 

critical  temperature,  165 

history,  20 

in  nature,  35 

in  the  air,  290 

preparation,  24-29 

physical  properties,  29 

separation  from  air,  23 

standard  for  atomic  weights,  351 

uses,  33 
Ozone,  144 

chemical  properties,  146,  147 

formation,  145-146 

heat  of  formation,  145 

history,  144 

in  the  air,  294 

physical  properties,  146 

preparation,  145 

uses,  148 


Paint-remover,  432 
Palladium,  625 
Palmitic  acid,  426 
Paper,  423 
PABACELSUS,  38,  414 


INDEX 


673 


Paraffin,  422 
Paris  green,  598 
PASCAL,  285 
PASTEUR,  285 
Peat,  196 

Perchlorates,  385,  386 
Perchloric  acid,  497 
Perchloric  anhydride,  386 
Periodic  law,  358,  388,  439 
Permanganates,  643 
Permanganic  acid,  643 
Permutit,  536 
Peroxides,  149,  150 
Persulphides,  280 
Persulphuric  acid,  280 
Petroleum,  422 
Pewter,  455 
Phase,  169 

Philosopher's  stone,  4 
Phlogiston,  52 
Phosgene,  177 
Phosphine,  405 
Phosphorescence,  400 
Phosphoric  acids,  401 
Phosphorus,  396 

action  with  bases,  405 

chemical  properties,  399 

discovery,  397 

iodides,  407 

occurrence,  397 

pentachloride,  406 

pentoxide,  401 

physical  properties,  398 

preparation,  397 

sulphides,  407 

trichloride,  406 

trioxide,  401 
Phosphorous  acid,  404 
Photography,  282,  692 
Picric  acid,  330 
Pitch-blende,  548 
Plants,  growth  of,  186 
Plaster  of  Paris,  539 
Platinum,  447,  623 

compounds,  624 
Poisons,  409 


Polymers,  265 
Polysulphides,  255 
Porcelain,  573 
Portland  cement,  575 
Potentials,  single,  485 
Potassium,  526 

carbonate,  527 

chlorate,  preparation,  384,  385 

chloride,  solubility,  506 

chloroplatinate,  529 

chromate,  629 

cobalticyanide,  619 

cyanide,  529 

dichromate,  629 

ferrate,  616 

ferricyanide,  616 

ferrite,  616 

ferrocyanide,  614 

hydroxide,  527 

manganate,  642 

metabisulphite,  264 

nitrate,  528 

heat  of  solution,  504 
solubility,  505 

permanganate,  642,  643 

sulphate,  solubility,  506 

tests  for,  529 
Pressure,  atmospheric,  79 

effect  on  boiling-point,  168 
equilibrium,  243 
freezing-point,  168 

electrolytic  solution,  484 

measurement,  78 
PRIESTLEY,  21,  22,  24,  88,  179,  256, 

284,  303,  314,  338 
Producer-gas,  202 
Propane,  421 
Properties,  physical,  chemical,  11 

variable,  characteristic,  9 
PROUST,  54 
Prussian  blue,  614 
Putty,  534 
Pyrolusite,  640 
Pyrometers,  475 
Pyrophosphoric  acid,  402 
Pyrosulphates,  276 


674 


INDEX 


Quartation,  599 
Quartz,  430 


Q 


R 


Radicals,  120,  131 
Radioactivity,  646,  648 
Radiometer,  8 
Radium,  648,  649,  651 
RAMSAY,  296,  299 
RAYLEIGH,  296 
Reactions,  endothermic,  141 
exothermic,  141 
of  double  decomposition,  221 
rate,  240 

reversible,  110,  234 
Realgar,  407,  411 
Reduction,  48,  632,  644 
Rhodium,  625 
RICHARDS,  356 
Rochelle  salt,  596 
RONTGEN,  652 
ROSE'S  metal,  455 
Rubidium,  515 
Ruby,  569 
Russia  iron,  464 
Rusting  of  iron,  617 
Ruthenium,  625 
RUTHERFORD,  285,  652,  655 

S 

Salt,  in  sea  water,  96 
Saltpeter,  Chile,  286 
Salts,  211 

acid  and  basic,  212 

complex,  510 

double,  510 

hydrolysis,  511 

ionization,  511 

properties,  503 

reaction  with  metals,  222,  465 

with  non-metals,  223 
solubility,  506 

in  acids,  507 

in  ammonia,  509 

in  other  salts,  509,  510 


Sandstone,  429 
Sapphire,  569 
SCHEELE,  100,  285 
SCHONBEIN,  144 
Science,  57 
Selenic  acid,  395 
Selenium,  392 

compounds,  393 
dioxide,  394 
Serpentine,  432 
Sewage,  97 
Siderite,  603 
Silica,  430 
Silicates,  430,  431,  466 

test  for,  433 
Silicic  acid,  430 
Silicon,  428,  429 
carbide,  192 
dioxide,  429 
hydrides,  433 
tetrachloride,  433 
tetrafluoride,  434 
Silk,  581 

artificial,  596 
Silver,  588 

chloride,  heat  of  formation,  143 
German,  455 
halides,  591 
nitrate,  591 
orthoarsenate,  404 
orthophosphate,  404 
oxide,  heat  of  formation,  142 
reaction  with  hydrogen  per- 
oxide, 153 
oxides,  590 
tests  for,  592 
Slag,  468,  606 
Smalt,  619 
Smaltite,  408 
Smithsonite,  555 
Soap,  426,  437,  535 

history,  4 
Soapstone,  432 
SODDY,  652 
Sodium,  516 

amalgam,  315 
arsenide,  410 


INDEX 


675 


Sodium  bicarbonate,  523 

carbonate,  520,  523 
hydrated,  93 

chloride,  516 

heat  of  solution,  505 
in  sea  water,  99 

dichromate,  630 

fluosilicate,  525 

hydroxide,  102,  618 

hypochlorite,  hydrolysis,  512 

iodate,  387 

metaborate,  437 

metaphosphate,  438 

nitrate,  321 

solubility,  506 

nitrite,  334 

oxide,  heat  of  formation,  142 

permanganate,  643 

peroxide,  28,  153 

persulphate,  280 

pyroantimonate,  416,  525 

salts  of,  525 

silicate,  432 

stearate,  426 

sulphate,  solubility,  506 

tests  for,  525 

tetrathionate,  280 

thiosulphate,  279 

zincate,  557 
Soldering,  438 
Solders,  452,  453 
Solids,  157 
Solubility  of  gases  in  liquids,  166 

of  salts,  504 

product,  502 
Solutions,  212 

boiling-point,  224 

freezing-point,  227 

normal,  229 
SOLVAY,  520,  522 
Specific  gravity,  159 

heat,  159 
Spectroscope,  525 
Spelter,  556 
Sphalerite,  408,  555 
Spinthariscope,  651 
Stannic  acid,  580 


Stannic  chloride,  581 
Stannous  chloride,  581 

hydroxide,  580 
Starch,  423,  525 

heat  of  combustion,  140 

as  test  for  iodine,  114 
Stearic  acid,  426 
Stearin,  426 
Steel,  607 
Steels,  alloy,  611 

high  speed,  612 
STEPHAN'S  law,  400 
Stereotype  metal,  455 
Stibine,  416 
Stone,  artificial,  432 
Strontium,  531,  645,  646 

hydroxide,  531 

sulphate,  solubility,  531 
Stucco,  540 
Sublimation,  167 
Substances,  pure,  15 
Substitution,  108,  131 
Sucrose,  424 
Sugar,  424 

heat  of  combustion,  178 
Sulpharsenites.     See  thioarsenites. 
Sulphates,  277 

test  for,  278 
Sulphides,  preparation,  253 

solubility  in  acids,  508 

test  for,  254 

use  in  analysis,  508 
Sulphites,  263 

test  for,  265 
Sulphur,  245 

chemical  properties,  248 

chlorides,  248 

extraction,  246 

occurrence,  245 

physical  properties,  247 

uses,  249 

Sulphur  dioxide,  catalytic  oxidation, 
270 

chemical  properties,  260 

critical  temperature,  165 

history,  256 

in  the  air,  294 


676 


INDEX 


Sulphur  dioxide,  occurrence,  256 
physical  properties,  260 
preparation,  257 
uses,  261,  262 

Sulphur  trioxide,  265 

Sulphuric  acid,  267 

action  on  metals,  461 
chemical  properties,  275 
contact  process,  269 
history,  267 

lead  chamber  process,  270 
manufacture,  268 
physical  properties,  274 
uses,  279 

Sulphurous  acid,  263 

Supercooling,  169 

Superheating,  169 

Sylvite,  526 

Symbols,  58 


Talc,  432 

Tartar,  cream  of,  524 

-emetic,  415 
Telluric  acid,  395 
Tellurium,  392 

chloride,  394 

dioxide,  394 

nitrate,  395 

sulphate,  394 
Tellurous  acid,  395 
Temperature,  critical,  164 

effect  on  equilibrium,  240 
Tempering  steel,  610 
Tenacity,  444 
Thermite,  568 
Thermochemistry,  140 
Thermometers,  77 
Theory,  57 

electrolytic  dissociation,  215 
THOMPSON,  J.  J.,  652 
Thioarsenates,  412 
Thioarsenites,  412 
Thiosulphates,  279 
Tin,  579.     See    also    stannous    and 
stannic. 

action  of  nitric  acid  on,  461 


Tin,  oxides,  579 
recovery,  114 
sulphides,  582 
symbol  for,  59 
test  for,  582 

T.  N.  T.,  330,  333 

Toluene,  199 

TOWNSEND  cell,  518 

Transition-points,  168 

Tripoli  powder,  430 

Tungsten,  636 
in  steel,  612 

Type  metal,  455 

U 

Ulexite,  436 
Uranium,  649 

compounds,  637 


Valence,  117 

effect  of  change  in,  440 

effect  on  chemical  properties,  626 

electronic,  389 

electrons,  653 

influence  on  basic  properties,  639 

of  metals  and  properties,  566,  578 

positive  and  negative,  390,  632, 

644 

VALENTINE,  BASIL,  414,  417 
VAN  HELMONT,  179 
VAN'T-HOFF,  242 
Vapor,  aqueous,  85 
Vaporization,  heat  of,  161 
Vegetables,  content  of  water,  89 
Venetian  red,  615 
Ventilation,  148,  292 
Verdigris,  598 
Vinegar,  211 

W 

WARD,  267 
War-gases,  177 
Water,  88 

action  with  metals,  40 

analysis,  336 


INDEX 


677 


Water,  chemical  properties,  91,  92 

composition,  94,  95 

critical  temperature,  165 

dissociation  by  heat,  113 

effect    of    pressure    on    boiling- 
point,  168 

effect   of   pressure   on   melting- 
point,  168 

electrolysis,  27,  39 

hard,  437,  535 

heat  of  formation,  142 
from  ions,  219 

heat  of  vaporization,  162 

history,  88 

ionization,  512 

occurrence,  88 

physical  properties,  89,  90 

purification,  96,  177 

sterilization,  148 

superheated,  169 

-gas,  201 

vapor,  90 

in  the  air,  292 
Waters,  natural,  96 


WATT,  473 
WELDON  mud,  642 
WELSBACH  gas-mantels,  401 
White  lead,  585 
Whiting,  534 

WlLHELMY,  239 

WOHLER,  567 
Wood,  196,  425 
WOOD'S  metal,  417,  455 


X 


Xenon,  301 
X-Rays,  646 


Zinc,  463,  465,  555 

action  of  carbon  dioxide  on,  183 

blende,  555 

chloride,  557 

hydroxide,  557 

oxide,  556 

salts,  557 

sulphide,  solubility  in  acids,  508 

tests  for,  557 


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tNGINEERII 

G  LIBRAE 

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FED  2  7  1947 

APR   1  1  K2CI. 

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889789 


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INTERNATIONAL  ATOMIC  WEIGHTS,  1921 


Sym- 
bol. 

Atomic 
Weight. 

Sym- 
bol. 

Atomic 
Weight. 

Aluminum  
Antimony            ...    . 

Al 
Sb 

27.1 
120.2 

Molybdenum  
Neodymium  

Mo 
Nd 

96.0 
144.3 

Argon 

A 

39  9 

Neon  

Ne 

20.2 

Ars6nic 

As 

74  96 

Nickel                      .    . 

Ni 

58.68 

Barium        

Ba 

137.37 

Niton    (radium   ema- 

Bismuth 

Bi 

208  0 

nation)  

Nt 

222.4 

Boron 

B 

10  9 

Nitrogen    .        

N 

14  008 

Bromine 

Br 

79  92 

Osmium 

Os 

190  9 

Cadmium 

Cd 

112  40 

Oxvcen  .  . 

o 

16.00 

CsBsium 

Cs 

132  81 

Palladium        

Pd 

106  7 

Calcium            

Ca 

40.07 

Phosphorus  

P 

31.04 

Carbon                      .  . 

c 

12  005 

Platinum  

Pt 

195.2 

Cerium 

Ce 

140  25 

Potassium 

K 

39  10 

Chlorine     

Cl 

35.46 

Praseodymium  

Pr 

140.9 

Chromium 

Cr 

52  0 

Radium             

Ra 

226.0 

Cobalt 

Co 

58  97 

Rhodium 

Rh 

102  9 

Columbium  

Cb 

93.1 

Rubidium  

Rb 

85.45 

Copper 

Cu 

63  57 

Ruthenium 

Ru 

101  7 

Dysprosium  

Dy 

162.5 

Samarium  

Sa 

150.4 

Erbium      .  .        .... 

Er 

167  7 

Scandium  

Sc 

45  1 

Europium  

En 

152.0 

Selenium  

Se 

79.2 

Fluorine           

F 

19.0 

Silicon  

Si 

28.3 

Gadolinum 

Gd 

157  3 

Silver         ..... 

Ag 

107  88 

Gallium 

Ga 

70  1 

Sodium 

Na 

23  00 

Germanium  

Ge 

72.5 

Strontium  

Sr 

87  63 

Glucinum 

Gl 

9  1 

Sulphur 

S 

32  06 

Gold.,  
Helium 

Au 
He 

197.2 
4  00 

Tantalum  
Tellurium  

Ta 
Te 

181.5 
127  5 

Holmium  

Ho 

163.5 

Terbium  

Tb 

159.2 

Hydrogen  

H 

1.008 

Thallium  

Tl 

204.0 

Indium 

In 

114  8 

Thorium  

Th 

232  15 

Iodine 

I 

126  92 

Thulium 

Tm 

168  5 

Iridium  

Ir 

193.1 

Tin  '  

Sn 

118  7 

Iron 

Fe 

55  84 

Titanium  .  .    . 

Ti 

48  1 

Kr}rpton  

Kr 

82.92 

Tungsten  

W 

184.0 

Lanthanum  

La 

139.0 

Uranium  

u 

238  2 

Lead                   

Pb 

207  .  20 

Vanadium  .    .  . 

V 

51  0 

Lithium 

Li 

6  94 

Xenon 

Xe 

130  2 

Lutecium  

Lu 

175.0 

Ytterbium  (Neoytter- 

Magnesium  

Mg 

24.32 

bium)  

Yb 

173  5 

Manganese 

Mn 

54  93 

Yttrium 

Yt 

89  33 

Mercury  

Hg 

200.6 

Zinc  

Zn 

65  37 

Zirconium  

Zr 

90.6 

